Transcript Document 7798739

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Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
The Puzzle of the Atom
 Protons and electrons are attracted to each
other because of opposite charges
 Electrically charged particles moving in a
curved path give off energy
 Despite these facts, atoms don’t collapse
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light.
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Types of electromagnetic radiation:
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Toupee?
The Wave-like Electron
The electron propagates
through space on an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.
Orbitals in outer energy levels DO penetrate into
lower energy levels. Penetration #1
This is a probability
Distribution for a
3s orbital.
What parts of the
diagram correspond
to “nodes” – regions
of zero probability?
The Great Niels Bohr
(1885 - 1962)
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
Bohr Model Energy Levels
Schrodinger Wave Equation

d

V 
8  m dx
h
2
2
2
2
 E
Equation for probability of a
single electron being found
along a single axis (x-axis)
Erwin Schrodinger
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
You can find out where the
electron is, but not where it
is going.
Werner
Heisenberg
OR…
You can find out where the
electron is going, but not
where it is!
Quantum Numbers
Each electron in an atom has a unique
set of 4 quantum numbers which describe
it.
 Principal quantum number
(n)
 Angular momentum quantum number (l)
 Magnetic quantum number (m)
 Spin quantum number
(s)
Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
The principal quantum number (n) cannot be
zero.
 n must be 1, 2, 3, etc.
Number of electrons
that can fit in a shell:
2n2
Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell) in
which the electron is located.
The angular momentum quantum number (l ) can be any integer
between 0 and n - 1.
l =3
f
An orbital is a region within an atom where there
is a probability of finding an electron. This is a
probability diagram for the s orbital in the first
energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
Magnetic Quantum Number
The magnetic quantum number, generally
symbolized by m, denotes the orientation of the
electron’s orbital with respect to the three axes in
space. The magnetic quantum number (ml) can be any integer
between -l and +l.
Pauli Exclusion Principle
No two electrons in an atom
can have the same four
quantum numbers.
Wolfgang
Pauli
Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin:
1

2
1

2
Assigning the Numbers
 The three quantum numbers (n, l, and m) are
integers.
 The principal quantum number (n) cannot be
zero.
 n must be 1, 2, 3, etc.
 The angular momentum quantum number (l )
can be any integer between 0 and n - 1.
 For n = 3, l can be either 0, 1, or 2.
 The magnetic quantum number (ml) can be any
integer between -l and +l.
 For l = 2, m can be either -2, -1, 0, +1, +2.
Principle, angular momentum, and magnetic
quantum numbers: n, l, and ml
Aufbau
1.
2.
3.
4.
5.
energy is emitted
energy is absorbed
no change in energy occurs
light is emitted
none of these
45%
36%
14%
5%
0%
it.
er
en
gy
is
..
em
er
en
gy
is
..
o.
s
ab
no
ch
ge
an
in
..
e.
i
ht
g
li
se
itt
...
m
ne
no
o
h
ft
e
es
ia
ra
d
re
d
w
in
fra
io
ra
d
ow
icr
...
av
es
s
av
e
s
9%
xra
ys
4%
0%
m
5.
17%
ra
y
4.
a
3.
m
2.
gamma rays
microwaves
radio waves
infrared radiation
x-rays
ga
m
1.
70%
1. 2
61%
2. 5
3. 10
17%
5. 6
13%
4%
6
18
10
2
4%
5
4. 18
½
–½
½
–½
½
14%
5%
5%
42
0½
0%
43
–2
–½
5.
36%
21
–1
½
4.
0
0
–1
–2
0
41%
30
0–
½
3.
1
0
1
3
2
s
11
0½
2.
1
3
2
4
4
m
s
1.
l
nlm
n
45%
3.
4.
5.
32%
14%
5%
5%
0%
10
13
s–
. ..
4.1
2x
10
5s
8.5
–1
0x
10
20
s–
9.1
. ..
2x
10
12
s–
. ..
3.2
0x
10
9s
–1
2.
3.00 x 1013 s–1
4.12 x 105 s–1
8.50 x 1020 s–1
9.12 x 1012 s–1
3.20 x 109 s–1
3.0
0x
1.
2.
3.
4.
5.
1s22s22p63s23p64s23d104p65s24d105p15d10
1s22s22p63s23p64s23d104d104p1
1s23s22p63s23p64s24d104p65s25d105p1
1s22s22p63s23p64s23d104p65s24d105p1
none of these
87%
9%
0%
0%
4%
1s
22
s2
2p
63
s2
3p
1s
.. .
22
s2
2p
63
s2
3p
1s
.. .
23
s2
2p
63
s2
3p
1s
.. .
22
s2
2p
63
s2
3p
.. .
no
ne
of
th
es
e
1.
Orbital filling table
Yet Another Way to Look at Ionization Energ
Element
Lithium
Configuration
notation
1s22s1
[He]2s1
____
1s
Beryllium
____
____
2p
____
____
2s
____
____
2p
____
[He]2s2p2
____
2s
____
____
2p
____
1s22s2p3
[He]2s2p3
____
2s
____
____
2p
____
1s22s2p4
[He]2s2p4
____
2s
____
____
2p
____
1s22s2p5
[He]2s2p5
____
1s
Neon
____
2s
1s22s2p2
____
1s
Fluorine
____
[He]2s2p1
____
1s
Oxygen
____
2p
1s22s2p1
____
1s
Nitrogen
____
[He]2s2
____
1s
Carbon
____
2s
1s22s2
____
1s
Boron
Noble gas
notation
Orbital notation
____
2s
____
____
2p
____
1s22s2p6
[He]2s2p6
____
1s
____
2s
____
____
2p
____
The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape
P orbital shape
There are three peanut-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.
d orbital shapes
Things get a bit more
complicated with the five d
orbitals that are found in
the d sublevels beginning
with n = 3. To remember
the shapes, think of:
“double peanut”
…and a “peanut
with a donut”!
Shape of f orbitals
Things get even more
complicated with the seven f
orbitals that are found in the
f sublevels beginning with n
= 4. To remember the
shapes, think of:
Flower
Electron configuration of the
elements of the first three series
Periodicity
}
Radius
Atomic Radius = half the distance between two nuclei
of a diatomic molecule.
 Influenced by three factors.
 Energy Level
 Higher energy level is further away.
 Charge on nucleus
 More charge pulls electrons in closer.
 Shielding
 Layers of electrons shield from nuclear pull.
 The electron on the outside
energy level has to look
through all the other energy
levels to see the nucleus
 The electron on the outside
energy level has to look
through all the other energy
levels to see the nucleus.
 A second electron has the
same shielding.
 Each atom has another
H
Li
energy level,
 So the atoms get bigger.
Na
 As we go down a group
K
Rb
 As you go across a period the radius gets smaller.
 Same energy level.
 More nuclear charge.
 Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
Table of
Atomic
Radii
 Cations form by losing electrons.
 Cations are smaller that the atom they come from.
 Metals form cations.
 Cations of representative elements have noble gas
configuration.
 Anions form by gaining electrons.
 Anions are bigger that the atom they come from.
 Nonmetals form anions.
 Anions of representative elements have noble gas
configuration.
Rb
K
Atomic Radius (nm)
Na
Li
Kr
Ar
H
Ne
10
Atomic Number
 The amount of energy required to completely remove
an electron from a gaseous atom.
 Removing one electron makes a +1 ion.
 The energy required is called the first ionization
energy.
 The second ionization energy is the energy required to
remove the second electron.
 Always greater than first IE.
 The third IE is the energy required to remove a third
electron.
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
 The greater the nuclear charge the greater IE.
 Distance from nucleus increases IE
 Filled and half filled orbitals have lower energy, so
achieving them is easier, lower IE.
 Shielding
 As you go down a group first IE
decreases because
 The electron is further away.
 More shielding.
 All the atoms in the same period have the same energy




level.
Same shielding.
Increasing nuclear charge
So IE generally increases from left to right.
Exceptions at full and 1/2 fill orbitals.
Ne
He
N F
l
First Ionization energy
l
H
C O
Be
Li
B
l
l
Na
Atomic number
Na has a lower
IE than Li
Both are s1
Na has more
shielding
Greater distance
Electron Affinity - the energy change
associated with the addition of an electron
 Affinity tends to increase across a period
 Affinity tends to decrease as you go down
in a group
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
 The tendency for an atom to attract electrons to itself
when it is chemically combined with another element.
 How fair it shares.
 Big electronegativity means it pulls the electron
toward it.
 Atoms with large negative electron affinity have larger
electronegativity.
 The further down a group the farther the electron is
away and the more electrons an atom has.
 More willing to share.
 Low electronegativity.
 Metals are at the left end.
 They let their electrons go easily
 Low electronegativity
 At the right end are the nonmetals.
 They want more electrons.
 Try to take them away.
 High electronegativity.
Ionization energy, electronegativity
Electron affinity INCREASE
Atomic size increases, shielding constant
Ionic size increases
Another Way to Look at Ionization Energy
Yet Another Way to Look at Ionization Energ
Summary of Periodic Trends
Put the following in order of
Decreasing atomic radius:
a) Cl,Ar,K
b) b) O, O-, O2c) Co, Rh, Ni
Now put them in order of Decreasing
ionization energy:
Prepare yourself to
^
C
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light.
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Types of electromagnetic radiation:
The energy (E ) of electromagnetic
radiation is directly proportional to the
frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
 = frequency, in units of hertz (hz, sec-1)
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Relating Frequency, Wavelength
and Energy
c 
E  h
Common re-arrangements:
E
hc

hc

E
Types of electromagnetic radiation:
PLANCK’S PRACTICE PROBLEMS
1. When we see light from a neon sign, we are observing
radiation from excited neon atoms. If this radiation has a
wavelength of 640 nm, what is the energy of the photon being
emitted?
2.
Light with a wavelength of 614.5 nm looks orange. What is the
energy, in joules, of a photon of this orange light?
3.
A photon of light produced by a surgical laser has an energy
of 3.027 x 10 -19 J. Calculate the frequency and the
wavelength of the photon.
 method that provides information on all the
occupied energy levels of an atom (that is, the
ionization energies of all electrons in the atom) is
known as photoelectron spectroscopy; this
method uses a photon (a packet of light energy) to
knock an electron out of an atom.
The photoelectron spectrum is a plot of the number of electrons emitted versus their
kinetic energy. In the diagram below, the “X” axis is labeled high to low energies so
that you think about the XY intersect as being the nucleus.


http://www.chem.arizona.edu/chemt/Flash/photoelectron.html
2p
6-
2- 1s
1-
2s
3p
3s
4s
Orbital names s, p, d,
and f stand for names
given to groups of lines
in the spectra of the
alkali metals. Early
chemists called the line
groups sharp, principal,
diffuse, and
fundamental.
Interpretations from the data:
1. There are no values on the y axis in the tables above. Using the Periodic Table
and Table 1, put numbers on the y axis.
2. Label each peak on the graphs above with s, p, d, or f to indicate the suborbital
they represent..
3. What is the total number of electrons in a neutral potassium atom?
 If a certain element being studied by an X-ray PES
displays an emission spectrum with 5 distinct kinetic
energies. What are all the possible elements that could
produce this spectrum?
 Determine the orbitals that the spectral lines are
originating from and then determine the elements
that have electrons in only these orbitals.