Download Chapter 4 – Arrangement of Electrons in Atoms

Chapter 4 – Arrangement of Electrons in Atoms
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Properties of Light
o electromagnetic radiation – all forms of energy that act as __________;
visible light, X-rays, ultraviolet, infrared light, microwaves, radio waves
o wavelength (λ) – the _________________ between corresponding points on
adjacent waves

For example, the visible light spectrum ranges from 400nm (violet/blue)
to 700nm (orange/red)
o frequency (ν) – the ______________ of waves that pass a given point in a
specific time, usually one second.
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One wave per second is referred to as 1 hertz (Hz)
Speed of Light
o c=λν
o c = speed of light = 3.0 x 108 m/s
o As wavelength increases, frequency _______________. This is an example of
an __________________ relationship.
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Energy States
o ______________ state – normal state; lowest energy state; where electrons
are typically located
o ______________ state – energized state; electrons are located farther away
from the nucleus than in their natural state
o “Nature is ___________!”

Matter will always find its _____________ energy state.

Atoms can be excited by adding ____________. When that energy is
used, the atom will return back to its ground state, releasing light (or
_______________).

The _______________ of light depends upon the atom.

Why do specific gases give off only specific frequencies of light?
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One interpretation – The Bohr Model
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Bohr Model
o To explain electrons within atoms, Niels Bohr designed a model of the atom.
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o Electrons can exist in these different orbits around the nucleus.
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These levels are like rungs on a _____________.

The further away an orbit is from the nucleus, the more ________ it has.

This model explains how excited electrons move to higher energy orbits.

As the electrons fall back down to their original energy level they
emit energy in the form of ___________.
o Shortcomings:

Bohr’s model only successfully explains the behavior of the _____ atom
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Bohr’s model can’t be applied to atoms with more than one _____ (due
to the complexity of the interactions between electrons and the nucleus)
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Photoelectric Effect
o photoelectric effect – the emission of ______________ from a metal (when a
light shines upon it)
o Einstein suggested that light has a dual ___________- _____________ nature.
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Each particle of light = _______________
o Why aren’t electrons emitted at all frequencies?
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They found that electrons were only emitted when a high enough
frequency of light hit the metal. There had to be a __________ frequency
or energy to pull or strip the electrons from the metal. Through this,
scientists like Einstein were able to see light’s dual wave-particle nature.
o Application - A solar powered calculator runs based on the frequencies of light it
receives.
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Quantum Mechanics
o Scientists, like Max Planck (a German physicist), concluded that energy is
found in packets or clusters, which are called ________________.
o All attempts to explain “quanta” and how it affects electrons and atoms are what
we consider “Quantum Mechanics.”
o And thus the field of Quantum Mechanics was born.
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Electrons as Waves
o After the photoelectric effect, things get a little weird for Quantum mechanics.
o Prior to the 20th century, scientists thought of matter as ______________.

They described atoms as billiard balls that can collide with each other
and bounce off obstacles.
o With quantum mechanics scientists had to acknowledge that small things (like
atoms and electrons) can also ___________ __________ ___________.
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Properties of Electrons and Small Particles
o Properties of Electrons and other small particles:

They are _____________. They have no fixed _______________.

They are able to spread out and be in ______ __________ ____
_______.
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Able to interfere with other _____________.
Back to Electrons as Waves
o Diffraction – the _____________ of a wave as it passes by the end of an
object
o Interference – overlapping waves that result in a ____________ of energy in
some areas and an ____________ of energy in others
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The Heisenberg Uncertainty Principle
o Quantum weirdness is the essence of the Heisenberg Uncertainty Principle.
o If electrons are both particles and waves, then where are they located in the
atom?
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Werber Heisenberg (German theoretical physicist) – 1927
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Electrons are detected by their interaction with __________. The reason
we see things is because of the photons striking off of the electrons on
that object. When the photon hits the electron, it changes the electron’s
position. Thus, there is always a basic uncertainty in trying to locate an
electron.
o In essence, HUP states that it is impossible to determine simultaneously both
the ______________ and ______________ of an electron.
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An electron is in multiple places at once until it is observed.
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The act of observing the electron locks it into place, but by observing it
the electron is changed and its velocity is different.
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Shrodinger’s Cat
o 1926 – Erwin Schrödinger (Austrian physicist)
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Shrodinger’s Cat – The cat is both ___________ and ___________.
Summary of Quantum Mechanics
o Since Quantum Mechanics changed scientist’s view of matter, it also changed
the understanding of atoms.
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1. Bohr’s model had set ______________ for electrons within the atom
(the orbits).
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2. Quantum mechanics claims that exact locations are
________________.
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3. The new result is an atom that has a nucleus surrounded by an
_____________ ___________.
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Within the cloud there is only a probability of finding the electron.
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The cloud still has energy levels like Bohr’s model.
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Within the energy levels there are differences that Bohr’s model
does not account for.
o Schrodinger also developed an ___________ that treated electrons as waves.
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Solutions to the equation = ________ ____________. These
wave functions determine the area where electrons orbit.
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Quantum theory – describes mathematically the wave properties
of electrons within atoms
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Fundamentals of the Quantum Atom
o Electrons are attracted to the __________ and will reside as close to the
nucleus as possible.
o Electrons ________ one another and orient themselves as far away as
possible.
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The combination of attraction to the nucleus and repulsion to other
electrons creates the unusual findings of this new model.
o An electron can exist in the same orbital with only one other electron, but both
must be traveling in opposite ______________ (aka. opposite “spins”).
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This is one reason the Bohr model is inadequate.
o Different _________ __________ will have different numbers and
arrangements of ______________.
o The energy of the ___________ determines how the electrons are arranged
within the energy level.
o In order to describe orbitals and the energy of the electrons within them,
scientists use ___________ _____________.
o Quantum numbers indicate the properties of the electrons and their orbitals.
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Quantum Numbers
o Principle Quantum Number – “n”, ______ __________ _______ occupied by
the electron
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As “n” increases, the electron’s energy and average distance from the
nucleus increase.
o Angular Momentum Quantum Number – “l”, indicates the ________ of the
orbital
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Exception: first main energy level
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Can be 0 and less than or equal to “n-1”
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Example: n = 2, l = 0 or 1
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L = zero, shape “s” (_____________)
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L = one, shape “p” (______________)
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L = two, shape “d” (__________ or __________)
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L = three, shape “f” (more complicated clovers and donuts)
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Example: 1s =
_______________________________________________
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Example: 2p =
_______________________________________________
o Magnetic Quantum Number – “m”, indicates the _______________ of an
orbital around the nucleus
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“s” sublevel – 1 orientation (m = 0)
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Spherical shape and thus only one possible orientation
“p” sublevel – 3 orientations (m = -1, 0, 1)
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Orientations x, y, and z (px, py, pz)
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“d” sublevel – 5 orientations (m = -2, -1, 0, 1, 2)
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“f” sublevel – 7 orientations (m = -3, -2, -1, 0, 1, 2, 3)
o Spin Quantum Number – has only two possible values (+ ½ and - ½), which
indicate the two fundamental spin states of an electron in an orbital
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Remember, 2 electrons per orbital (maximum) – both with ___________
spins.
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A Tour of the Quantum Atom
o 1st Energy Level
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Simplest
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Closest to nucleus
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Only _____ orbital (a spherical “s” orbital)
o 2nd Energy Level
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Further from nucleus
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More complicated orbitals
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_____ s orbital (spherical) and _____ p orbitals (dumbbell)
o 3rd Energy Level
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Even further away from the nucleus
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More complicated orbitals
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_____ s orbital (spherical), _____ p orbitals (dumbbell), _____ d
orbitals
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Four of the d orbitals are clover-leaf shaped, the fifth is looks like
a dumbbell with a donut around its middle
o 4th and 5th Energy Levels
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Same number of orbitals, but 5th level is further from nucleus and higher
in energy
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_____ s orbital, _____ p orbitals, _____ d orbitals, _____ f orbitals
o 6th Energy Level
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_____ s orbital, _____ p orbitals, _____ d orbitals
o 7th Energy Level
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_____ s orbital, _____ p orbitals
Atom Analogy
o Think of the atom as a city.
o The city has 7 neighborhoods (7 energy levels)
o Each neighborhood can have up to four types of streets (s, p, d, and f
sublevels)
o Each street can have up to 7 houses (maximum of 7 orbitals)
o Each house can hold 2 electrons (single orbitals)
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Orbital Math Tips
o When counting the number of electrons in an energy level, sublevel, or orbital,
remember:
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Up to 2 electrons per orbital
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Sublevels contain all orbitals for that shape:
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“s” sublevel = 1 orbital, “p” sublevel = 3 orbitals, “d” sublevel = 5 orbitals,
“f” sublevel = 7 orbitals
o An energy levels can have multiple sublevels
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Orbital Math Rules
o _____________ Principle – Nature is Lazy! An electron will fill the lowest
energy orbit possible.
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The order of the orbitals is as follows:
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1s →2s→ 2p→ 3s→ 3p→ 4s→ 3d→ 4p→ 5s→ 4d→ 5p→ 6s→ 4f→
5d→ 6p→ 7s→ 5f→ 6d→ 7p
o __________ Rule – if equal energy orbitals are present, electrons will fill them
separately before pairing up.
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Only applies to p, d, and f sublevels (because they have multiple
orientations)
o _________ _____________ Principle – no two electrons will have the same
four quantum numbers
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Even if you have two electrons in the same energy level, same sublevel,
and same orbital – they will have different spins
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Notation
o 3 methods to represent the location of electrons within an atom:
o (1) Orbital Notation
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Represents each electron as an arrow pointing up or down.
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The direction of the arrow indicates spin.
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Examples = Atomic numbers 1-10, 23, 46, 73
o (2) Electron Configuration Notation
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Abbreviated form of orbital notation that uses superscripts to represent
electrons.
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Examples = Atomic number 3, 5, 13, 46, 80, 108
o (3) Noble Gas Notation
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Shortest of the three types
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Lists the noble gas from the period above the element, and electron
configuration to represent all remaining electrons.
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*Leave up “Let’s Practice” slides for students to refer to. Do the first three
examples, and then have the students do the rest.*
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Examples: Atomic numbers = 3, 5, 13, 46, 80, 108