Transcript Ch. 5 Electrons in Atoms

Ch. 5 Electrons in Atoms
This Quiz = Begin learning/memorizing
the elements and symbols for
1-20,22,24-30, 33,35,36,38
Next Quiz = Learn
47,50,51,53,54,78-80,82,92
Ch. 5 Electrons in Atoms
KC #1: Rutherford/Bohr Model was too
simplistic and couldn’t explain the huge
range of varied properties found in
atoms. Extreme changes in properties
meant complex structure
Fluorine #9 = most reactive element and a gas
Neon #10 = 2nd most stable/non-reactive
element and a gas
Sodium #11 = one of the most reactive elements
and a solid
Ch. 5 Electrons in Atoms
Bohr Model expanded on Rutherford by
proposing that energy in atoms comes in
“increments” and electrons are found on
Energy Levels
Bohr Model of the Atom
In 1914 Niels Bohr proposed that the
energy levels for the electrons in an
atom are quantized
En = -hcRH (1/n)2
En = (-2.18  10-18 J)(1/n2)
Where n = 1, 2, 3, 4, …
n=1
n=2
n=3
n=4
Ch. 5 Electrons in Atoms
•Light and energy absorption and emission
•Same idea, just with particle accelerators and
electron microscopes
•The specific energies that an electron can have
•The amount of energy absorbed to move an
electron from one energy level to another
•No, in fact they are different for every atom,
regardless of energy levels
•Worked well for explaining the behavior of
hydrogen, but not other elements (F, Ne, Na)
Ch. 5 Electrons in Atoms
•A set of mathematical equations that described the
range of possible energies an electron could have
•Determines the allowed energies an electron can
have and how likely it is to find the electron in
various locations around the nucleus
•A mathematically defined region based on a %
likelihood of find an electron with the given area
Ch. 5 Electrons in Atoms
•The energy levels around an atom (1-7)
•The divided regions of space within an Energy
level (spdf)
•Each energy sublevel corresponds to an orbital
(probability cloud) of a different shape (all
‘s’ are alike)
•s p d f, shape and amounts of energy
•s=2 p=6 d=10 f=14 use the formula above and
the periodic table
Ch. 5 Electrons in Atoms
Let’s stop for a minute, and review….
1. Electrons are found around the nucleus in an
extremely complex organization
2. This organization is predicted and described by
a set of math equations called Schroedinger’s
Equations.
3. The complex organization of the electrons
explains and predicts the properties and behavior
of elements
E. L.
Sublevels
1
s
2
s
p
3
s
p
4
spdf
5
s p d f “g”
6
7
Orbitals
d
Electrons
1
2
1 3
2 6
1 3 5
2 6 10
1 3 5 7
2 6 10 14
Noble Gas Shorthand
Orbital
Diagrams
Ch. 5 Electrons in Atoms
- The arrangement of electrons of an atom in its ground state (no
absorbed energy) into levels and orbitals
- Some elements “break the rules” like Cu and Cr on page 136
Give an example of an electron configuration.
1s2 2s2 2p6 3s2 3p6
Give an example of a Noble Gas shorthand configuration.
[Ne] 3s2 3p5
Given an example of an orbital diagram.
_ _
  _
1s
2s
2p
Ch. 5 Electrons in Atoms
Ch. 5.2
- Aufbau, Pauli, Hund
Ch. 5.3
- Energy waves that travel at the speed of light but
have differing wavelength and frequency
- The range or spectrum of electromagnetic
radiation as they gradually change wavelength and
frequency
λ = wavelength
a = amplitude
ƒ = frequency (waves/s)
v = speed of light
300,000,000 m/s
Electromagnetic Spectrum
Light as a Wave
- Inversely proportional
- They would “tie” as they all
travel the speed of light
c = ln
l = wavelength (m)
n = frequency (s-1)
c = speed of light
(3.00  108 m/s)
Ch. 5 Electrons in Atoms
- When atoms absorb energy, electrons move into
higher energy levels. These electrons lose energy by
emitting light when they return to lower energy levels
- The lowest possible energy of an electron and the
energy level it moves to after absorbing energy
- The light emitted by an electron moving from a
higher to a lower energy level has a frequency directly
proportional to the energy change of the electrons
Blackbody Radiation & Max Planck
The classical laws of physics do
not explain the distribution of
light emitted from hot objects.
Max Planck solved the problem
mathematically (in 1900) by
assuming that the light can only
be released in “chunks” of a
discrete size (quantized like
currency or the notes on a piano). l = wavelength (m)
We can think of these “chunks” as n = frequency (s-1)
particles of light called photons. h = Planck’s constant
E = hn
E = hc/l
(6.626  10-34 J-s)
Line Spectrum of Hydrogen
n=6 n=5
n=4
In 1885 Johann Balmer, a Swiss
schoolteacher noticed that the
frequencies of the four lines of
the H spectrum obeyed the
following relationship:
n = k [(1/2)2 – (1/n)2]
Where k is a constant and
n = 3, 4, 5 or 6.
n=3
Ch. 5 Electrons in Atoms
- The energies emitted by an Hydrogen atom that fall in the
visible portion of the spectrum
- Atomic emission spectra are the patterns formed when the
unique shade of light generated by a certain element passes
thru a prism and is separated into different freq it contains
- Again, Bohr could explain Hydrogen with a single
electron, but not the complexities of more electrons
- As a particle.
- Almost exclusively as a wave
Ch. 5 Electrons in Atoms
- That energy came in quantums or “packets”
- A “particle” of energy
- Waves and particles were seemingly opposites,
yet light exhibits both characteristics
Photoelectric Effect
In 1905 Albert Einstein
explained the
photoelectric effect using
Planck’s idea of quantized
photons of light. He later
won the Nobel Prize in
physics for this work.
Ch. 5 Electrons in Atoms
- Experiment performed in multiple ways, but
classically as red and violet light shined on sodium
metal
- The red light never “worked” to eject electrons and
the violet light always “worked”
- Proved that energy has particle characteristics
-The red light did not provide the “right” amount of
energy to make the electrons move an energy level
Louis DeBroglie & the WaveParticle Duality of Matter
While working on his
PhD thesis (at the
Sorbonne in Paris)
Louis DeBroglie
proposed that matter
could also behave
simultaneously as an
particle and a wave.
l = h/mv
l = wavelength (m)
v = velocity (m/s)
h = Planck’s constant
(6.626  10-34 J-s)
Only important for matter that has a very small mass
(electrons). As mass goes up wavelength goes down. We
will see later in some ways electrons behave like waves.
Electron Diffraction
Electron Diffraction
Pattern
Transmission Electron Microscope
Heisenberg & the Uncertainty Principle
While working as a
postdoctoral
assistant with Niels
Bohr, Werner
Heisenberg
formulated the
uncertainty principle.
We can never precisely
know the location and
the momentum (or
velocity or energy) of
an object. This is only
important for very
small objects.
Dx Dp = h/4p
Dx = position uncertainty
Dp = momentum uncertainty
(p = mv)
h = Planck’s constant
- Classical Mechanics adequately
describes the motions of larger
bodies, while quantum
mechanics describes the motions
of subatomic particles and atoms
and waves
Ch. 5 Electrons in Atoms
- Einstein theorized that if we ever perfectly froze
an atom, it would violate Heisenberg, so it would
instead form a an entire atom probability cloud
- Uncertainty coupled with duality of light and
matter led to Schroedinger’s equations
Schrodinger and Electron Wave
Functions
In Schrodinger’s wave mechanics
the electron is described by a
wave function, Y. The exact
wave-function for each electron
depends upon four variables,
called quantum numbers they are
Erwin Schrodinger, an
Austrian physicist,
proposed that we think of
the electrons more as
waves than particles.
This led to the field
called quantum mechanics.
n = principle quantum number
l = azimuthal quantum number
ml = magnetic quantum number
ms = spin quantum number
Ch. 5 Electrons in Atoms
- These equations led to the idea of “quantum
numbers” or what we generically refer to as
electron configuration.
- Further study of this concept through an
experiment called the “Double Slit Experiment”
determined that electrons also exhibited wavelike
properties. This led to the Idea of Quantum
Mechanics and the theory or energy absorption and
emission.
Ch. 5 Electrons in Atoms
1. At atom’s electrons actually absorb the energy
2.The electrons absorb a photon of just the right amount of
energy
3.The just right amount of energy is called a quantum
4.The electron is on an energy level called its ground state
5.It absorbs the photon containing a quantum of energy, then
jumps to its excited state
6.The electron is now energized and unstable
7.The electron releases the same amount of energy and goes
back to its ground state
8.The energy is released typically in the form of visible light or
color
9.Repeats
Y2 = Probability density
s-orbital Electron Density
(where does the electron
spend it’s time)
# of radial nodes = n – l – 1
Velocity is proportional
to length of streak,
position is uncertain.
Position is fairly certain,
but velocity is uncertain.
Schrodinger’s quantum mechanical picture of the atom
1. The energy levels of the electrons are well known
2. We have some idea of where the electron might be at a
given moment
3. We have no information at all about the path or trajectory
of the electrons
s & p orbitals
d orbitals
# of nodal planes = l
Electrons produce a magnetic field.
All electrons produce a magnetic
field of the same magnitude
Its polarity can either be + or -,
like the two ends of a bar magnet
Thus the spin of
an electron can
only take
quantized values
(ms=+½,-½),
giving rise to the
4th quantum
number
Single Electron Atom
Multi Electron Atom