Download 5.1 Revising the Atomic Model

5.1 Revising the Atomic Model >
Chapter 5
Electrons In Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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5.1 Revising the Atomic Model >
CHEMISTRY
& YOU
Why do scientists use mathematical
models to describe the position of
electrons in atoms?
Shown here is a lifesized model of a skier,
but not all models are
physical. In fact, the
current model of the
atom is a mathematical
model.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
Objective
Energy Levels in Atoms
Understand how the atomic
model was revised.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
• It explained only a few simple
properties of atoms.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
• It explained only a few simple
properties of atoms.
• It could not explain the chemical
properties of elements.
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5.1 Revising the Atomic Model > Energy Levels in Atoms
Limitations of Rutherford’s Atomic Model
• It explained only a few simple
properties of atoms.
• It could not explain the chemical
properties of elements.
For example, Rutherford’s model
could not explain why an object
such as the iron scroll shown here
first glows dull red, then yellow,
and then white when heated to
higher and higher temperatures.
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5.1 Revising the Atomic Model >
 1913, Niels Bohr develops a new
atomic model
 Bohr stated that the electrons orbit
the nucleus like the planets orbit the
sun.
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5.1 Revising the Atomic Model >
Each possible electron orbit in
Bohr’s model has a fixed energy.
•The fixed energies an electron can
have are called energy levels.
•Each energy level further from the
nucleus is of greater energy
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5.1 Revising the Atomic Model >
 There are 7 different energy levels
• Each energy level can contain a different
amount of electrons
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5.1 Revising the Atomic Model >
Electrons orbit the
nucleus in circular
paths of fixed energy
(energy levels).
n=1  first energy level
n=2 second energy lev
n=3  third energy level
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Niels Bohr’s
Model (1913)
5.1 Revising the Atomic Model >
Highest energy
level for carbon
is n = 2 (2 rings).
Valence electrons
– electrons in the
outermost energy
level
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5.1 Revising the Atomic Model >
Bohr's model:
-electrons orbit the nucleus like planets
orbit the sun
-each orbit can hold a specific maximum
number of electrons
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5.1 Revising the Atomic Model >
The Rutherford model could not
explain why elements that have been
heated to higher temperatures give off
different colors of light.
The Bohr model explains how the
energy levels of electrons in an atom
change when the atom emits light.
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5.1 Revising the Atomic Model >
Energy and Atoms
Ground State: the lowest energy state of an atom.
- An electron absorbs energy (photon) and moves
from the ground state to an excited state.
Excited State: when an atom contains excess energy
(has higher potential energy).
When an excited atom returns to ground state it gives
off light (the energy it has gained as
electromagnetic radiation).
Example: Neon signs
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5.1 Revising the Atomic Model >
Absorption
An electron
absorbs energy
(photon) and
moves from the
ground state to
an excited
state.
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E
4
E
3
E
2
E
1
5.1 Revising
Atomic
Model come
>
Whatthe
goes
up…must
down!
Emission
When an electron in the excited state returns
to the ground state it emits a photon.
E
4
E
3
E
2
E
= h =E -E
photon
3 1
E
1
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5.1 Revising the Atomic Model >
Absorption and Emission
Emitting photons creates light
or electromagnetic radiation
Electromagnetic radiation in
the visible light spectrum has
color!
These photons have
wavelengths that correspond
to their color.
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5.1 Revising the Atomic Model >
Unfortunately, Bohr’s model only
applied to hydrogen atoms and did not
apply to other atoms.
That led scientists to question his
model
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5.1 Revising the Atomic Model > Wave Mechanical Model
Today, the modern description of
electrons in atoms is called the Quantum
Mechanical Model.
• The wave model tells you the
probability of finding an electron in an
atom (the exact path of an electron is
not known)
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5.1 Revising the Atomic Model >
Do Now
What does the Bohr model say about
electrons?
Why is this model incorrect?
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5.1 Revising the Atomic Model >
 There are 7 different energy levels
• Each energy level can contain a
different amount of electrons
• There are 4 different types of sublevels
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5.1 Revising the Atomic Model >
Energy levels (n=1, n=2 ….)
Sublevels (s,p,d,f)
orbitals
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5.1 Revising the Atomic Model >
The sublevels each can contain a different
amount of electrons
• s – 2 electrons
• p – 6 electrons
• d – 10 electrons
• f – 14 electrons
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5.1 Revising the Atomic Model >
• An atomic orbital is a region of space in
which there is a high probability of finding
an electron.
• These orbitals have different shapes
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5.1 Revising the Atomic Model >
Energy level 1 has only an s sublevel – total
of 2 eEnergy level 2 has the s and p sublevels –
total of 8 eEnergy level 3 has the s, p, and d sublevels –
total of 18 eEnergy level 4 has the s, p, d, and f sublevels
– total of 32 e-
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5.1 Revising the Atomic Model >
Do Now
1. How many principal energy levels are there? ____
2. What are the four sublevels? _____
3. How many orbitals does the f sublevel hold? _____
4. How many electrons can each orbital hold? _____
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5.1 Revising the Atomic Model >
Do Now Answers
1. How many principal energy levels are there? 7
2. What are the four sublevels? s, p, d , f
3. How many orbitals does the f sublevel hold?
4. How many electrons can each orbital hold?
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7
2
(n=7)
5.1 Revising the Atomic Model >
Chapter 5
Electrons In Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement
in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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5.2
5.1 Electron
Revisingarrangement
the Atomic Model >Electron Configurations
Aufbau Principle
6p
5d
6s
4f
5p
4d
Increasing energy
5s
4p
4s
3p
3s
2p
2s
1s
29
3d
According to the aufbau principle,
electrons occupy the orbitals of
lowest energy first. In the aufbau
diagram, each box represents an
atomic orbital.
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5.2
5.1 Electron
Revisingarrangement
the Atomic Model >Electron Configurations
Hund’s Rule
According to Hund’s rule, electrons
occupy orbitals of the same energy in a
way that makes the number of electrons
with the same spin direction as large as
possible.
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5.2
5.1 Electron
Revisingarrangement
the Atomic Model >Electron Configurations
Hund’s Rule
Three electrons would occupy three
orbitals of equal energy as follows.
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5.2
5.1 Electron
Revisingarrangement
the Atomic Model >Electron Configurations
Hund’s Rule
Three electrons would occupy three
orbitals of equal energy as follows.
Electrons then occupy each orbital
so that their spins are paired with the
first electron in the orbital.
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Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.
5.2
Electronthe
arrangement
5.1 Revising
Atomic Model >
Electron configuration describes the
placement of the electrons
• Example: Hydrogen: 1s1
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5.1 Electron
Revisingarrangement
the Atomic Model >Electron Configurations
5.2
Look at the orbital filling diagram of the oxygen atom.
• An oxygen
atom contains
eight
electrons.
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Electron Configurations of Selected Elements
Element
1s
2s
2px 2py 2pz
3s
Electron
configuration
H
1s1
He
1s2
Li
1s22s1
C
1s22s22p2
N
1s22s22p3
O
1s22s22p4
F
1s22s22p5
Ne
1s22s22p6
Na
1s22s22p63s1
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5.1 Revising the Atomic Model >
Chapter 5
Electrons In Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra
and the Quantum
Mechanical Model
35
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5.3
Emission
Spectra
5.1 Atomic
Revising
the Atomic
Model >
What gives gas-filled lights their colors?
An electric current
passing through the gas
in each glass tube
makes the gas glow
with its own
characteristic color.
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5.1 Revising the Atomic Model >
5.3 Electromagnetic Radiation and Energy
By the year 1900, there was enough experimental
evidence to convince scientists that light consisted of
waves.
The wavelength, represented by (the Greek letter
lambda), is the distance between the crests.
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5.1 Revising the Atomic Model >
5.3 Electromagnetic Radiation and Energy
The frequency, represented by (the Greek letter
nu), is the number of wave cycles to pass a given
point per unit of time.
The SI unit of waves per second is called the hertz
(Hz).
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5.1 Revising the Atomic Model >
The frequency ( ) and wavelength ( ) of light are
inversely proportional to each other.
As the wavelength increases, the frequency
decreases.
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5.1 Revising the Atomic Model >
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5.1 Revising the Atomic Model >
5.3 Electromagnetic Radiation and Energy
According to the wave model, light consists of
electromagnetic waves.
Electromagnetic radiation - a form of energy that
exhibits wavelike behavior as it travels through
space.
– All electromagnetic radiation travels at the
speed of light:
c = 3.0 X108 m/s
Electromagnetic radiation includes radio waves,
microwaves, infrared waves, visible light, ultraviolet
waves, X-rays, and gamma rays.
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5.1 Revising the Atomic Model >
Rutherfo
https://www.youtube.com/watch?v=cfXzwh3KadE
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5.1 Revising the Atomic Model >
Wave Description of Light
Equation relating frequency and
wavelength:
c = 
c = speed of light (m/s)
 = wavelength (m)
 = frequency (Hz or s-1)
=c

=c

c is constant, so is , so as frequency
increases, wavelength decreases (inversely
proportional).
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Light as a Wave: Problems
5.1 Revising the Atomic Model >
c = 
1) If c = 3.00 x 108 m/s and  = 1 x 1019s-1 ,
what does  equal?
2) What is the frequency of light () if its
wavelength () is 4.34 X 10-7 m?
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5.1 Revising the Atomic Model >
Visible light of different wavelengths can be
separated into a spectrum of colors.
In the visible spectrum, red light has the longest
wavelength and the lowest frequency.
Violet light has the shortest wavelength and the
highest frequency.
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5.1 Revising the Atomic Model >
Atomic Emission Spectrum
Atomic Emission Spectrum- a beam of light
separated into a series of specific frequencies
(and therefore specific wavelengths) of visible
light.
– produced when electrons fall back to ground
state
– the energy emitted in the fall give off specific
patterns (colors) of light
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The Hydrogen-Atom Line
5.1 Revising the Atomic Model >
Emission
Spectrum
The Hydrogen-Atomic Emission
Spectrum
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5.1 Revising the Atomic Model >
The Hydrogen-Atomic Emission Spectrum
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5.1 Revising the Atomic Model >
Atomic Emission Spectrum of Na, He, Ne, and Mercury
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5.1 Revising the Atomic Model >
A fluorescent lamp or a fluorescent tube is a low
pressure mercury-vapor gas-discharge lamp that
uses fluorescence to produce visible light. An
electric current in the gas excites mercury vapor
which produces short-wave ultraviolet light that then
causes a phosphor coating on the inside of the bulb
to glow.
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5.1 Revising the Atomic Model >
The Quantization of Energy
Planck’s constant (h)
= 6.63 x 10-34 Js
Energy
E = h
Frequency ()
- German physicist Max Planck (1858–1947)
showed mathematically that the amount of
radiant energy (E) of a single quantum
absorbed or emitted by a body is proportional
to the frequency of radiation ().
- h is Planck’s constant = 6.63 x 10-34 Js
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5.1
RevisingProblems
the Atomic Model >
Energy
E = h
h = 6.63 x 10-34 Js
1. If the frequency () = 1.15 x 1012 s-1 ,
what is the energy of the radiation?
1. What is the energy of a photon of
microwave radiation with a frequency of
3.20 × 1011s-1?
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5.1 Revising the Atomic Model >
Photons
Einstein proposed that light could be
described as quanta of energy that
behave as if they were particles.
These light quanta are called photons.
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5.1 Revising the Atomic Model > Key Concepts
Bohr proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
The quantum mechanical model determines
the allowed energies an electron can have and
how likely it is to find the electron in various
locations around the nucleus of an atom.
Each energy sublevel corresponds to one or
more orbitals of different shapes, which
describe where the electron is likely to be
found.
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5.1 Revising the Atomic Model > Glossary Terms
• energy level: the specific energies an
electron in an atom or other system can
have
• quantum: the amount of energy needed to
move an electron from one energy level to
another
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5.1 Revising the Atomic Model > Glossary Terms
• quantum mechanical model: the modern
description, primarily mathematical, of the
behavior of electrons in atoms
• atomic orbital: a mathematical expression
describing the probability of finding an
electron at various locations; usually
represented by the region of space around
the nucleus where there is a high probability
of finding an electron
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P =
#e- =
#N =
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = atomic #
=9
#e- =
#N =
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = # P
=9
#N =
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = atomic mass - # P
= 10
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
draw the nucleus with
protons & neutrons
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9P
10N
5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
9P
10N
#N = 10
how many electrons can fit in the first orbit?
62
5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
9P
10N
#N = 10
how many electrons can fit in the first orbit?
2
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left?
64
9P
10N
5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
how many electrons are left?
7
65
9P
10N
Bohr's
Model
of the
5.1 Revising
the Atomic
Model
> Atom
e.g. fluorine:
#P = 9
#e- = 9
9P
10N
#N = 10
how many electrons are left?
7
how many electrons fit in the second orbit?
66
Bohr's
Model
of the
5.1 Revising
the Atomic
Model
> Atom
e.g. fluorine:
#P = 9
#e- = 9
9P
10N
#N = 10
how many electrons are left?
7
how many electrons fit in the second orbit?
8
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5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
68
9P
10N
5.1 Revising the Atomic Model >
Bohr's Model of the Atom
e.g. fluorine:
#P = 9
#e- = 9
#N = 10
How many valence electrons?
7
69
9P
10N