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Transcript
The Atom and Its
Properties
Chapter 4 – Nucleus
Chapter 5 – Electron
Configuration
Chapter 4 Objectives
• Describe
an atom’s structure and
differentiate among the particles that
make it up.
• Identify the numbers associated with
elements and explain their meaning .
• Realize that the number of protons in a
nucleus defines an element.
• Calculate the average atomic mass
given isotopes and relative abundance
2
Chapter 4 Vocabulary
Chapter 4.3
Chapter 4.1
• Atomic Number
• Dalton’s Atomic
• Isotope
Theory
• Mass Number
• Atom
• AMU (Atomic Mass Unit)
Chapter 4.2
• Average Atomic Mass
• Electron
• Nucleus
• Proton
• Neutron
3
Modern View of the Atom
The nucleus is where the protons and neutrons
are located and contain most of the atom’s
mass.
4
Protons, Neutrons and Electrons
5
Particle
Symbol
Charge
Electron
Proton
Neutron
ep+
n0
11+
0
Relative
Mass
1/1840
1
~1
•Sometimes Atomic Symbols are
Displayed as:
6
Isotope Examples
Isotope Atomic Mass
Number Number Number
Number Number Protons Neutron Electron
Cr-52
24
52
33
222
7
42
86
What’s all this amu business?
•
•
•
8
To simplify a system of indicating atomic
masses since protons and neutrons have such
extremely small masses, scientists have
assigned the carbon-12 atom a mass of
exactly 12 atomic mass units. (amu)
The mass of 1 amu (1/12 the mass of carbon12) is very nearly equal to the mass of a single
proton or neutron but not the same.
1 amu = 1.66 x 10-24 grams
•Isotopes and Mass Number
• Your
text (p. 119) shows how to
calculate the mass number for Cl given
the % abundance of the isotopes.
• Let’s do this for another element: Li
• 6Li
• 7Li
is 7.59 % abundant; 6.015 amu
is 92.41% abundant; 7.015 amu
• Method
1: Use percentages. Think of
this as a sample of 100 atoms.
9
•In Tabular Form
Species
Mass
(amu)
6Li
6.015
(isotope)
7Li
7.015
(isotope)
Li
(100 atoms)
Li (atom)
10
Abundance Mass x
Abundance
%
(Weighted
share)
7.59
45.65
92.41
648.26
100.00
694.41
6.94
Average Atomic Mass
• The
average mass of an atom is found by
weighting the natural abundances of its
isotopes.
• Lithium
• 6Li
• 7Li
(Method 2): Change % to fraction.
6.015 amu 7.59% = 0.0759
7.015 amu 92.41% = 0.9241
Mass (amu) Frac abund Mass share
Avg mass = 6.015 amu x 0.0759 = 0.46 amu
7.015 amu x 0.9241 = 6.48 amu
6.94 amu/atom
11
•Displayed on Periodic Table
12
13
Electrons in Atoms
Chapter 5
•Chapter 5 Objectives
Compare wave and particle matters of
light
• See how frequency of light emitted by an
atom is unique to that atom
• Compare and contrast the Bohr and
quantum mechanical models of the atom
• Express the arrangements of electrons in
atoms through orbital notations, electron
configurations, and electron dot structures
•
15
Chapter 5 Vocabulary
Chapter 5.1
• Electromagnetic Radiation
• Wavelength
• Frequency
• Amplitude
• Electromagnetic Spectrum
• Quantum
• Photoelectric Effect
• Photon
• Atomic Emission Spectrum
16
Chapter 5.2
• Ground State
• Quantum Number
• Quantum Mechanical
Model of the Atom
• Atomic Orbital
• Principal Quantum
Number
• Principal Energy
Level
• Energy Sublevel
•Wave Nature of Light
• Electromagnetic
radiation displays
wavelike behavior as it travels through
space
• Waves can be described by several
common characteristics
17
•Characteristics of a Wave
•
•
Waves transfer energy
Properties of waves:
• Frequency (ν –
pronounced ‘nu’) Number of vibrations
per unit time – Hz
(cycle/second)
• Wavelength (λ) Distance between
points on two
consecutive waves
• Speed of wave is
Frequency x
wavelength
Speed = ν x λ
Frequency is the
number of waves that
hit this point in one
second.
Amplitude
λ
1.5
1
0.5
0
0
200
400
600
-0.5
-1
-1.5
18
800
1000
1200
•Electromagnetic Spectrum
Note: All EM Radiation travels at 3.00 x 108 m/s
19
•Electromagnetic Spectrum
The speed of light (3.00  108 m/s) is the
product of it’s wavelength and frequency
c = λν.
20
The Electromagnetic Spectrum
– all light is energy
21
•Electromagnetic Spectrum
• Gamma
Rays – Highest frequency,
shortest wavelength. Can pass through
most substances
• X Rays – Lower frequency than Gamma
rays. Can pass through soft body tissue
but can’t pass through bone.
• Ultraviolet (UV) Rays – Part of sunlight
that causes sunburn
22
•Electromagnetic Spectrum
• Visible
Light – Sensitive to our eyes.
Allows us to see color
• Infrared – Less energy and longer
wavelength than visible light. Felt as
heat given off a heater or near a fire
• Radio Waves – Lowest frequencies on
the EM spectrum. Used by radio and
over-the-air TV.
23
An electromagnetic wave has a frequency of 6.0 x 104 Hz.
Convert this frequency into its corresponding wavelength.
Which region of the EM spectrum does this correspond
to?
lxn=c
l = c/n
l = 3.00 x 108 m/s
6.0 x 104 /s
l = 5.0 x 103 m
It’s a radio wave
(~103 meters)
24
7.1
•Practice
Problems
What is the
frequency of green
light, which has a
wavelength of
520 nm.
• A radio station
broadcasts at
94.7 MHz. What is
the wavelength of
the broadcast?
•
25
Answers
•Nature of Light
• Max Planck (18581947) studied the
different light
emitted from heated
objects
• Matter can only gain
or lose energy in
small specific
amounts
26
•Nature of Light
• A quantum is the minimum amount of
energy that can be gained or lost by an
atom
• The energy of EM radiation is proportional
to its frequency (E α ν)
27
•Photons
Albert Einstein (18791955) proposed that
while a beam of light
had wavelike
characteristics, it also
can be thought of as
a stream of tiny
particles (or bundles of
energy) called
photons
• Each photon carries a
quantum of energy
•
28
•Particle Nature of Light
The photoelectric effect is when electrons are
emitted from a metal’s surface when light
of a certain frequency shines on it.
29
When copper is bombarded with high-energy electrons,
X rays are emitted. Calculate the energy (in Joules) of
the X-rays if their frequency is 1.95 x 1018 Hz.
E=hxn
E = 6.63 x 10-34 (J•s) x 1.95 x 1018 /s
E = 1.29 x 10 -15 J
30
7.2
Ch. 5.2 – Quantum Theory of the Atom
Each element has only certain specific
frequencies of light that are emitted
when atoms absorb energy and become
excited
Where do we see this?
fireworks
neon signs
stars
31
•Hydrogen Spectrum
32
•Balmer Plot
•
In 1885, Johann Balmer observed the lines
of the spectrum fit this surprisingly simple
formula:
 1 1 
 RH  2  2 
l
 n1 n2 
1
•
Where n1 =2 and n2 = 3, 4, 5, etc.
33
1/Labmda, m-1
•Balmer Plot
2500000
2400000
y = 1.0972E+07x + 4.0238E+02
2300000
R² = 1.0000E+00
2200000
2100000
2000000
1900000
1800000
1700000
1600000
1500000
0.12
0.14
0.16
0.18
0.2
0.22
0.24
1/2^2 - 1/n^2
RH is the slope of this line, 1.0972 x 107 m-1
34
•Electronic Energy Transitions
• Neils
Bohr (18851962) proposed
the model the
hydrogen atom
(1913) to explain
the discreet
nature of the
hydrogen
spectrum.
35
•Electronic Energy Transitions
• Neils
Bohr’s model the atom (1913)
Electrons exist only in discrete, “allowable”
energy levels
• Energy is involved in moving electrons
from one energy level to another
• Principal quantum number (n) - specifies
the electron’s major energy level
• The lowest energy is n=1, the next lowest in
n=2, etc.
•
36
•Bohr’s Model of the Atom (cont’d)
Bohr suggested that an electron moves around the
nucleus in only certain allowed circular orbits.
n=2
n=1
37
•Energy Absorption/Emission
38
39
•Atomic Emission Spectra
40
•Origin of Line Spectra
Balmer series
41
42
•Quantum or Wave
Mechanics
Schrodinger applied idea of ebehaving as a wave to the
problem of electrons in atoms.
He developed the WAVE
EQUATION
E. Schrodinger
1887-1961
43
Solution gives set of math
expressions called WAVE
FUNCTIONS.
Treated electrons as wavelike
particles that became the
Quantum Mechanical Model of
the Atom.
Waves
•Wave motion: wave length and nodes
“Quantization” in a standing wave
44
•Hydrogen Atom Solution
Where:
a0 is the Bohr Radius given by a0 = 4πεoh2/me2
Generalized Laguerre Polynomial
m here is quantum number
Constant = 2.18 x 10-18 J
m is mass of electron
45
46
•Atomic Orbitals-Hydrogen
47
•Orbitals
No more than 2 e- assigned to an orbital
• Orbitals grouped in s, p, d (and f)
sublevels
•
s orbitals
p orbitals
d orbitals
48
s orbitals
p orbitals
d orbitals
s orbitals
No.
orbs.
p orbitals
d orbitals
1
3
5
2
6
10
No. e-
49
Energy Levels and
Sublevels
• Sublevels
are grouped in
energy level.
• Each energy level has a
number called the PRINCIPAL
QUANTUM NUMBER, n which
indicates relative size and
energy of the orbitals
• Row on PT indicates n
50
•Energy Levels and
Sublevels
n=1
n=2
n=3
n=4
51
•QUANTUM NUMBERS
The shape, size, and energy of each
orbital is a function of quantum
numbers:
n (principal)
l (angular)
• Note:
 Energy Level
 sublevel (s, p, d, or
f) which is its shape
There are other quantum numbers
that we will NOT discuss in detail. The ‘n’
and ‘l’ are sufficient.
52
QUANTUM NUMBERS
Symbol
Values
n (principal)
1, 2, 3, ..
l (angular)
More commonly
noted as:
53
Description
Orbital size
and energy level
0, 1, 2, 3, … n-1
Orbital shape
s, p, d, f, …n-1
or type
(energy sublevel)
Levels and Sublevels
When n = 1, then l = 0
Therefore, in n = 1, there is just 1 type
of sublevel
and that sublevel has a single orbital
This sublevel is labeled s (“ess”)
Each level has 1 orbital labeled s, and it
is SPHERICAL in shape.
54
•Types of
Atomic
Orbitals
55
Types of Atomic Orbitals
56
•s Orbitals— Always Spherical
Dot picture of
electron cloud
in 1s orbital.
Surface density
4πr2y versus
distance
See Active Figure 6.13
57
Surface of 90%
probability sphere
•p Orbitals
The three p orbitals lie 90o apart in space
58
•2px Orbital
59
3px Orbital
•Hydrogen-like Orbitals
(at most two electrons/orbital)
n
1
2
3
4
60
Sublevel
(l)
s
s
p
s
p
d
s
p
d
f
Orbitals
Max.
Orbital
Max
Elec
n2
2n2
s
1
2
s
px, py, pz
4
8
s
px, py, pz
dxy,dxz,dyz,dx2-y2, dz2
9
18
s
px, py, pz
dxy,dxz,dyz,dx2-y2, dz2
And 7 f orbitals
16
32
Chapter 5.3 Vocabulary
•
•
•
•
•
•
61
Electron configuration
Aufbau principle
Pauli Exclusion Principle
Hund’s Rule
Valence Electron
Electron Dot Structure
•Electron Configurations
• An
atom’s electron configuration is
the arrangements of electrons in
the atom.
• Electrons are arranged to minimize
energy.
• In other words, electrons fill up the
lowest energies possible first. This is
the Aufbau Principle.
62
•Assigning Electrons to Atoms
• Electrons
generally assigned to
orbitals of successively higher energy.
• For
• For
H atoms, E depends only on n.
many-electron atoms, energy
depends on both n and l.
63
•Energy Level Diagram of
Hydrogen
64
•Assigning Electrons to Subshells
In H atom all subshells of
same n have same energy.
In many-electron atom:
a) subshells increase in
energy as value of n + l
increases.
b) for subshells of same n +
l, subshell with lower n is
lower in energy.
•
65
Orbitals and Their Energies
Many-Electron Atoms
66
•Aufbau Diagram -- Filling Electron
Orbitals
1s 2s
3s
4s
5s
6s 7s
3p
4p
5p 6p
3d
4d 5d
8s
Start here
n+l=1
n+l=2
n+l=3
2p
n+l=4
n+l=5
n+l=6
4f
n+l=7
6d
7p
Haven’t gotten this
far. What orbitals
are being filled with
elements 110-118?
5f
n+l=8
The orbital with the lower ‘n’ is lower in energy if the n+l number is the same.
67
•Writing Atomic Electron Configurations
Two ways of writing
configs. One is
called the spdf
notation.
spdf notation
for H, atomic number = 1
1
1s
value of n
68
no. of
electrons
value of l
•Pauli Exclusion Principle
No two electrons in the
same orbital can have
the same spin. One
electron is spin up, the
other is spin down.
69
•Writing Atomic Electron
Configurations
Two ways of
writing configs.
Other is called
the orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
It would be a violation of the Pauli
exclusion principle to have both of
these electrons as spin up or both
as spin down.
70
•Electron Configurations
and the Periodic Table
71
•Lithium
Group 1 (1A)
Atomic number = 3
1s22s1  3 total electrons
3p
3s
2p
2s
1s
72
Interactive Periodic
Table
Ground State Electron Configurations
73
•Carbon
Group 14 (4A)
Atomic number = 6
1s2 2s2 2p2 
6 total electrons
3p
3s
2p
2s
1s
74
Here we see for the first time
HUND’S RULE. When
placing electrons in a set of
orbitals having the same
energy, we place them singly as
long as possible.
Electron Configuration for Elements 11-18
Noble gas notation uses noble gas symbols in
brackets to shorten inner electron configurations
of other elements.
75
Sodium
Group 1 (1A)
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses noble gas notation)
Note that we have begun a new
period.
All Group 1A elements have
[core]ns1 configurations.
76
Electron Configurations
and the Periodic Table
77
Transition Metals
All 4th period elements have the
configuration [argon] nsx (n - 1)dy
and so are d-block elements.
Chromium
78
Iron
Copper
79
Electrons in Energy Levels
• Electrons
fill up levels from lowest
energy to highest energy (Aufbau
Principle)
• Outermost
electrons are called
Valence Electrons.
• When atoms come close together it is
the Valence Electrons that interact.
80
Valence Electrons
• How
to determine which
electrons are in outer shell?
• Write
down electron
configuration of an element in
noble-gas configuration
• Whatever electrons are
displayed in the highest energy
shell (n) only are valence
electrons (Main Group elements)
81
Lewis Dot Diagrams
• How
do we represent Valence
Electrons?
• By a Lewis Dot Diagram
• Rules for Lewis Dot Diagrams:
• Use
the Elemental Symbol
• Use 1 dot to represent each valence
electron
• The symbol represents the nucleus and all
the inner (core) electrons.
• Examples:
. ..
..
•
Li·, Be: , ·C·,
. ·Cl:,
.. :Ne:
..
82
Valence Electrons
Examples
• O given by [He]2s22p4 so it has 6 valence
electrons.
•
•
83
Ga given by [Ar]3d104s24p1 has 3 valence
electrons.
•Electron Dot Representation
84