Transcript Electrons!

Electrons!
The nuclei of atoms (protons
and neutrons) are NOT
involved in chemical reactions,
BUT ELECTRONS ARE!
The first clue that early scientists
had to the existence of electrons
was ATOMIC SPECTRA.
ATOMIC SPECTRA
Continuous Spectrum – If white light is
passed through a prism, a continuous
spectrum (rainbow) can be observed.
The order of the colors is ROY G BIV
red, orange, yellow, green, blue, indigo, violet
Red has the longest wavelength and the lowest
energy.
Violet has the shortest wavelength and the highest
energy.
Electromagnetic Spectrum

The Atomic Spectrum produces visible
light and is part of the Electromagnetic
Spectrum.
Electromagnetic Radiation

Visible light is one form of electromagnetic
radiation. Some other examples are:
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Gamma rays – highest energy
X-rays
Ultraviolet rays
Visible light
Infrared
Microwaves
Radiowaves
Electricity – lowest energy
Atomic Spectra, continued
Bright Line Spectrum – Compounds and
metallic ions produce characteristic colors
when placed in a flame. When these
“flame tests” are passed through a prism,
a bright line spectrum (black background
with colored slits) can be seen.
Bright Line Spectrum
A spectrum is produced when radiation
from light sources is separated into
different wavelength components.
 The bright line spectrum contains radiation
of only specific wavelengths of light.

Hydrogen example
Spectral lines are the
fingerprints of the
elements!
Atomic Spectra, continued
Spectral Tubes – by passing an electrical
current through a sealed tube containing a
gas of an element, a bright line spectrum
can be produced.
The Modern Theory of Light
 Light
has dual properties:
Light has both wave and particle properties
 Light is packets of energy called quanta or
photons
 The amount of energy in each quantum
depends on the color of the light
 Violet light has the most energy.

Equations to calculate the energy
in light
Wave properties: f = c / λ
where:
f = frequency (s)
c = speed of light
(or 3.0 x 108 m/s)
λ = wavelength (m)
Equations to calculate the energy
in light
Particle properties:
where
E=hf
E = Energy (J)
h = Planck’s constant
(6.63 x 10-34 J · s)
f = frequency (s)
Bohr’s Model

Since energy must be absorbed for an eto move to a higher level and that energy
is emitted when it jumps back to ground
state, the total amount of energy can be
determined by:
∆E = Ef – Ei = Ephoton = hv
 Bohr states that only specific frequencies of
light that satisfy this equation can be
absorbed or emitted.
Think About It!

As an electron in a hydrogen atom jumps
from the n=3 orbit to the n=7 orbit, does it
absorb or emit energy?
ABSORB
ENERGY!
Calculating light energy, cont.

Since frequency is proportional to both
wavelength and energy, there is a
relationship between the wavelength of
light and energy. This is the idea of
“quantized energy” – a specific color of
light that represents a specific energy or
quantum of energy or photon.
Ground State vs. Excited State
The lowest energy state of an electron is
called “ground state.”
 When an electron is in a higher energy
state, it is called “excited state.”

Ground State

An electron in its lowest possible energy
level
Ground state e- configuration:
12Mg
2–8–2
Each electron energy level contains the
maximum number of electrons that it can
hold.
Excited State

An electron that has gained energy and
moved to a higher energy level – it’s very
unstable!
Excited state e- configuration:
12Mg
1–8–3
2–7–3
2–6–4
Hydrogen’s Line Emission Spectrum
n=6
n=5
n=4
n=3
n=2
e
n
e
r
g
y
410 Wavelength
(nm)
434
486
656
n=1
UV wavelengths
Energy in Level Jumping

The energy that an electron absorbs or
emits as it jumps can be calculated using
the following equation:
∆E = Ef – Ei = Ephoton = hv

Using the Energy State constant, which is
equal to E = (-2.18 x 10-18 J) (1/n2), it is
possible to calculate the difference in
energy.
Energy Calculations

∆E = hv = hc/λ = (-2.18 x
10-18
1
1
J) ----- - ----n 2f
n 2i
Example: If the electron moves from
ni = 3 to nf = 1, the calculation is:
(-2.18 x 10-18 J)
= -1.94 x 10-18 J
1
1
----- - ----12
32
8
9
Knowing the energy of the emitted photon,
we can calculate either its frequency or its
wavelength.
c
hc
 Wavelength: λ = --- = ---v
∆E

(6.626 x 10-34J-s) (3.00 x 108 m/s)
-1.94 x 10-18 J
= 1.03 x 10-7 m
Quantum Mechanics

Where does this energy come from?
Quantum mechanics is a field of physics that
answers this.
 Electrons absorb a specific number of
photons of energy when they are excited
(heated or absorb some other form of
energy). The electrons are not stable in that
state and emit photons of energy (in the form
of light or other forms of electromagnetic
energy) when they return to normal states.

Electrons can act like waves and
particles (DeBroglie, 1924.) He based
this on the Bohr’s original model and
used it to explain why electrons don’t
fall into the nucleus.
There are areas of an electron cloud
where the electrons are most dense,
and therefore, you would be most likely
to find an electron.
DeBroglie’s Theory

He suggested that as the e- moves around
the nucleus it is associated with a
particular wavelength, and that this
wavelength depends on its mass (m) and
on its velocity (v) (where h = Planck’s
constant.)
h
λ = -----mv
DeBroglie, continued
mv is the momentum.
 DeBroglie also used the term matter
waves to describe the wave characteristics
of material particles.
 This works only because the mass of an
electron is so small.

Heisenberg’s Uncertainty
Principle

The dual nature of matter limits how
precisely we can know the exact location
and momentum of any electron because of
its very small mass.
Schrodinger’s Wave Equation
His equation incorporates both the
wavelike behavior and the particle-like
behavior of an electron, and involves
complex calculus.
 His work yields a set of wave functions
and corresponding energies. These
functions are called orbitals.

Orbitals and Quantum
Numbers
Orbitals
Each orbital describes a specific
distribution of electron density in space.
 Each orbital has a characteristic energy
and shape.
 The lowest energy orbital in Hydrogen has
an energy of -2.18 x 10-18 J and a specific
shape as seen in the upcoming slides.

Models
Bohr model had a single quantum number,
n, to describe an orbit.
 The wave mechanical model uses three
quantum numbers, n, l, and m to describe
the orbitals.

Wave Mechanical Model Orbitals
The first quantum number, the principal
quantum number, n, can have positive
integral values of 1, 2, 3, etc.
 As n increases, the orbital becomes larger
and the electron is farther from the nucleus.
 An increase in n also means that the electron
has a higher energy and is less tightly bound
to the nucleus.

Second Quantum Number
The second quantum number, the angular
momentum quantum number, l, can have
integral values from 0 to (n-1) for each
value of n.
 This quantum number defines the shape
of the orbital.
 The value of l is generally designated by
the letters, s, p, d and f, corresponding to l
values of 0, 1, 2 and 3 respectively.

Third Quantum Number
The third quantum number, the magnetic
quantum number, ml, can have integral
values between -l and l, including 0.
 This quantum number describes the
orientation in space.
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Quantum
Number
Symbol
Principal
n
Angular
Momentum
Magnetic
Spin
l
ml
ms
Possible Values 1, 2, 3 . . . ∞
0, 1, 2, 3, …n - 1
s, p, d, f
-1. . .0. . .+1
±½
Electron/Orbital
Characteristic
Shape
Orientation
Magnetic
Spin
Size
Electron Shells
The collection of orbitals with the same
value is called an electron shell. All
orbitals that have n = 3, are said to be in
the third shell.
 The set of orbitals that have the same n
and l values is called a subshell.

Orbital Shapes
s Orbitals
The electron density is spherically
symmetric.
 The electron density at a given distance
from the nucleus is going to be the same
no matter what direction from the nucleus.
 The most probable distance to find an
electron in the s orbital, 0.529 Å, is
identical to the orbit predicted by Bohr.

p Orbitals
The electron density is not spherical.
 It is dumbbell-shaped and concentrated in
2 lobes separated by a node at the
nucleus.
 Starting with the n = 2 shell, each shell has
3 p orbitals in different spatial orientations
(-1, 0, +1)

d and f Orbitals
When n = 3 or greater, there are d orbitals.
 In each shell there are 5 possible values
for ml. (-2, -1, 0, 1, 2)

Principle n
Angular
Momentum l
Magnetic ml
Spin ms
Sublevel
designation
Number of
Orbitals
1
0
0
±½
1s
1
2
0
1
0
-1, 0, +1
±½
±½
2s
2p
1
3
3
0
1
2
0
-1, 0, +1
-2, -1, 0, +1, +2
±½
±½
±½
3s
3p
3d
1
3
5
4
0
1
2
3
0
-1, 0, +1
-2, -1, 0, +1, +2
-3, -2, -1, 0, +1, +2, +3
±½
±½
±½
±½
4s
4p
4d
4f
1
3
5
7
n=4
I
n
c
r
e
a
s
i
n
g
e
n
e
r
g
y
n=3
n=2
4f
7 x 2e- =14
4d
32
5 x 2e- = 10
4p
3 x 2e- = 6
3d
5 x 2e- = 10
4s
1 x 2e- = 2
18
3p
3 x 2e- = 6
3s
1 x 2e- = 2
2p
3 x 2e- = 6
2s
1 orbital x 2e- = 2
8
n=1
1s
1 orbital x 2e- = 2
Many Electron Atoms
The shapes of the orbitals for manyelectron atoms are the same, but it greatly
changes the energies of the orbitals.
 In a many-electron atom, the electronelectron repulsion cause the different
subshells to be at different energies.
 For a given value of n, the energy of an
orbital increases with a value of l.

Electron Spin
Electron spin is a property that states that
each electron is like a tiny sphere spinning
on its own axis.
 This leads to a new quantum number, ms,
the spin magnetic quantum number.
 The Pauli Exclusion Principle states that
no two electrons in an atom can have
the same set of 4 quantum numbers.

Electron Configurations
An electron configuration is a distribution
of electrons among the various orbitals of
an atom.
 Paramagnetic – refers to an atom having
one or more unpaired electrons. (Ex: Li, B,
C, N, Na)
 Diamagnetic – means that all the
electrons in an atom are paired.
