Download Chapter 4 Arrangement of Electrons in Atoms

Chapter 4
Arrangement of Electrons in Atoms
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Section 4.1
The Development of a New Atomic Model
Objectives:
Discuss the significance of the photoelectric
effect and the line-emission spectrum of
hydrogen to the development of the atomic
model.
 Describe the Bohr model of the hydrogen
atom.

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Electromagnetic radiation = energy that exhibits
wavelike behavior
All electromagnetic radiation travels at the same speed.
Different types are the result of different wavelengths
and frequencies.
This spectrum shows the different forms of
electromagnetic radiation.
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
Light is a small part of the
electromagnetic spectrum.
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Absorption/Emission Spectra

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


When light strikes a metal, the metal ejects efrom the surface and creates an electric current
= Photoelectric effect
Remember the cathode-ray tube?
When the glass tube was filled with a pure gas
and an electric current passed through, the gas
will gave off light.
Different gases give off different colors of light.
If this light is passed through a prism, a series
of bright lines is seen (emission spectrum).
Every element has a distinct emission spectrum.
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So, what’s happening to create
the emission spectrum?
Carbon
Oxygen
Iron
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Rutherford’s model of the atom
provided information about the
structure of atoms, it did not explain
where the electrons were located in
the space surrounding the nucleus.
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Max Planck
In 1900, Max Planck suggested that objects could
give off energy in small, specific amounts he called
quanta.
 A quantum (singular of quanta) is the minimum
amount of energy that can be lost or gained by an
atom.

Albert Einstein

Einstein proposed that different elements
require different frequencies of energy to eject
electrons
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So scientists agreed…
Light is a form of energy
Different colors of light have different levels of
energy on the electromagnetic spectrum
Atoms of different elements had different values
for a quantum (the minimum amount of energy
they can gain and lose)
When atoms of different elements absorb their
quantum of energy they can temporarily eject
electrons
Different colors of light are created by the
different levels of energy being absorbed, and then
given off, by electrons.
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The question still remained:
Why would different elements absorb
different amounts of energy and then
give off different light colors when
energy was applied and they ejected
electrons?
What was it about their structures
that allowed this?
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Neils Bohr Model of the H Atom
When an e- is hit by light energy, it absorbs the energy.
 If the energy is of the correct frequency (quantum), the e- will
jump to another energy level (excited state vs ground state).
 The electron cannot stay in excited state so it falls back to the
ground state. It cannot take the energy with it so it releases
the energy in the form of LIGHT energy!


We see this energy as different colors of light!
When the electron
returns to ground
state, it emits energy
in the form of
different colors of light –
depending on frequency
of energy gained.
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Section 4.2
The Quantum Model of the Atom
Objectives:
Compare and contrast the Bohr model and the
quantum model of the atom.
 List the 4 quantum numbers, and describe their
significance.
 Relate the number of sublevels corresponding
to each of an atom’s main energy levels, the
number of orbitals per sublevel, and the number
of orbitals per main energy level.

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Today’s Quantum Model of the Atom

Orbital -- a 3-d region around the nucleus that
indicates the probable location of an e- .

Quantum theory -- describes mathematically
the wave properties of e- and other very small
particles

Quantum numbers -- numbers or letters that
specify the properties of atomic orbitals and the
properties of e- in orbitals
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4 Quantum Numbers
1.
Principal Quantum Number
2.
Angular Momentum Quantum Number
3.
Magnetic Quantum Number
4.
Spin Quantum Number
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Principal quantum number (n) - indicates
the main energy level occupied by the en = 1,2,3,4, etc.
Representation of Bohr’s proposal of orbitals. Lower
numbered levels are closer to nucleus and of lower energy.
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Angular Momentum Quantum Number –
indicates the orbital and its shape = s, p, d, f
The maximum number of electrons in an energy
level is 2n2, thus a shell with n = 2 may hold a
maximum of 8 electrons.
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Magnetic Quantum Number – orientation of orbital around
the nucleus (the axis, or axes, it is on or between);
uses x, y, z, xy, yz, etc.
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These orbitals can overlap, cause
interference with each other, and affect
over all energy of each other.
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Spin Quantum Number – spin state of an
electron in an orbital – can be +½, -½ often represented with arrows
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Review
4 quantum numbers describe each e- .
1.
Principal quantum number - indicates the main energy level
occupied by the e- = 1,2,3,4, etc.
2.
Angular Momentum Quantum Number – indicates the orbital
and its shape = s, p, d, f
3.
Magnetic Quantum Number – orientation of orbital around
the nucleus (the axis, or axes, it is on or between) = x, y, z, xy,
yz, etc.
4.
Spin Quantum Number – spin state of an electron in an orbital
– can be +½, -½ - often represented with
So…to accurately describe the energy (location) of electron, all 4
quantum numbers must be used.
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Section 4.3
Electron Configurations
Objectives:
List the total number of electrons needed to
fully occupy each main energy level
 State and explain the Aufbau principle, the Pauli
exclusion principle, and Hund’s rule.
 Write orbital notations, electron-configuration
notation and noble-gas notation for atoms.

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Electron Configuration
Electron configuration -- the arrangement of ein an atom
 Electron configurations summarize the locations
of each e- in atoms.
 Like all systems in nature, e- tend to assume
arrangements that have the lowest possible
energies
 Ground-state configuration -- lowest energy
arrangement of the e- for each element

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3 basic rules help govern these
ground-state configurations.
1.
2.
3.
Aufbau principle -- an e- occupies the lowest-energy
orbital possible
Pauli exclusion principle -- no 2 e- in the same atom
can have the same set of 4 quantum numbers
Hund’s rule -- orbitals of equal energy are each
occupied by one e- before any orbital is occupied by
a second e- , and all e- in singly occupied orbitals
must have the same spin
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Types of Electron Configurations
1. Orbital Notation – includes all 4 quantum
numbers (spin is indicated with an arrow)
Fluorine – 9 e-
_____
1s
_____
2s
_____
2px
_____
2py
_____
2pz
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Complete Orbital Notations for
Elements 1 – 10 on the Periodic Table
1.Hydrogen
2.Helium
3.Lithium
4.Beryllium
5.Boron
6.Carbon
7.Nitrogen
8.Oxygen
9.Fluorine
10.Neon
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Types of Electron Configurations
2. Electron Configuration Notation – includes
principal and Angular Momentum (combines
magnetic and ignores spin)
Chromium
24 e-
1s2 2s2 2p6 3s2 3p6 4s2 3d4
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Complete Electron Configurations for
Elements 11 – 20 on the Periodic
Table
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Sodium
Magnesium
Aluminum
Si
P
S
Cl
Ar
K
Ca
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Types of Electron Configurations
3. Noble Gas Notation – abbreviates part of the
electron configuration by using the Noble
Gas Symbol just prior to the element and
adds the rest of the electron configuration
Magnesium 12eMg
=
[Ne] 3s2
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Complete Noble-Gas Notations for
the following elements:
S
Ca
Rb
Li
O
Al
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Chapter 4 Study Guide
1.
2.
3.
4.
5.
6.
7.
8.
Know the contributions of Plank and Einstein to the Quantum
Theory.
Be able to define “ground state” and “excited state” and know
what happens to cause an atom to moved from ground state to
excited state and what happens when it moves back to the
ground state.
Know the 4 quantum numbers by name and descriptions.
Know the 3 rules: Aufbau, Pauli Exclusion, and Hund and know
how they are applied when doing configurations.
Know the order of energy levels and orbitals from lowest energy
to highest energy.
Be able to write configurations: orbital notation, electron
configuration, and noble gas notations.
Be able to use configurations to identify elements.
Know the s, p, d, and f blocks on the Periodic Table so you can
double check your work.
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Class work
1.A quantum of electromagnetic energy is called
a(n)
.
2.A bright-light spectrum of an atom is caused by
the energy released when electrons
.
3. Explain the difference in an atom’s ground state
and excited state.
4.Name and describe the information provided
by all four quantum numbers.
5.Describe the shapes of the s, p, and d orbitals.
6.How many possible orientations can each
of the following sublevels have?
p sublevel?
d sublevel?
f sublevel?
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7. How many electrons can be held in a
p sublevel?
d sublevel?
f sublevel?
8. How many total electrons can be held in each
energy level?
1st –
2nd –
3rd –
4th –
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9. Explain each of the following rules governing
electron configurations.
Aufbau principle –
Pauli exclusion principle –
Hund’s rule –
10. Do the following for an atom of Cesium.
Orbital Notation
Electron Configuration
Noble Gas Notation
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