Transcript Chemistry Ch 4

Chemistry Ch 4
Arrangement of Electrons in
Atoms
Rutherford’s Model
Gold Foil Experiment
 Discovered the nucleus
 Did not explain where the electrons were
in an atom
 Why were they not attracted to the protons
in the nucleus?

Background Info
Electromagnetic Radiation are types of
energy.
 We describe these as waves. Only a
portion of these waves are visible to us
(the visible light waves).
 Each type of wave has different wave
characteristics: frequency, wavelength,
and the amount of energy it contains.

Relationship between light and
electrons

The electromagnetic spectrum includes all
types of electromagnetic radiation that
behaves as waves.
Light can behave as a wave (as in the
spectrum) and as a particle of matter like a
marble.
 Speed of all electromagnetic radiation is
3.0 X 108 m/s.

The Electromagnetic Spectrum
Wave Characteristics
Wavelength (λ) -the distance between two
points on a wave (measured in nm)
 Frequency (v) –the number of waves that
pass in a given point in one second
 The speed of light (c) – is a constant
 C= λ v
 Wavelength and Frequency are inverse of
each other (opposite).

Converting Wavelength Units
Wavelength is measured in nm.
 Speed is measured in m/s.
 They must both have the same unit so we
must convert nm to m to use it in the
equation.
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You broke your big toe! The x ray they take of toe uses
waves that have a length
4.0 X 10-7m. ( 1 meter = 1 X 109 nm)
What is the wavelength in nm? (l = 400 nm)
Frequency and Wavelength
Problems

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Calculate n for a l = 700 nm.
(Red 700nm
n = 4.3 x 1014 /s
or 4.3 x 1014 Hz)
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A purple light has a frequency of 7.42 x
1014 Hz.
What is its wavelength in nm?

(l = 404 nm)
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Lets look at a video to see what we are
going to be learning and what the
scientists were investigating…

Video of electron behavior as waves and particles
Photoelectric Effect
Experiments by Einstein and others in the
1900s tried to explain the interactions
between light and matter that were not
explained with the wave theory
 Their research led them to discover the
dual wave particle nature. How
electromagnetic radiation behaves as
waves and as particles.

The evidence for this was the
Photoelectric Effect experiment, which
explained how light, that is usually thought
of as a wave, can also behave like a
particle of matter.
 Lois de Broglie wondered if electrons
(matter), normally thought of as a particle,
maybe have some wave properties too.

Photo of Photoelectric Effect

The wave theory predicted that light of any
frequency could supply enough energy to
eject an electron from its position.

However, no electrons were emitted if a
light’s frequency was below a certain
minimum, regardless of how long the light
was shone.
Max Planck suggested that an object
emits energy in small, specific amounts,
called quanta.
 A quantum is the minimum quantity of
energy that can be lost or gained by an
atom.
 E= hv
 h = Planck’s constant 6.626 X 10-34 Js

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Photons are the “particles of light” that
carry a certain amount of energy.

The energy of a photon depends on the
frequency of the wave.
In order for an electron to be ejected from
a metal surface, the electron must be
struck by a single photon with the
minimum energy required to knock the
electron lose. (supports the particle theory)
 Because E = hv, the minimum energy
needed corresponds to the frequency
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Energy Problem

What is the energy of a photon whose
frequency is 3.0 X 1012 Hz?
E = hv
Where h= 6.626 x 10-34 J/Hz
E =[6.626 x 3.0] 10 (-34+12) J
1.99 X 10-21 J
An observation was that different metals
required different amounts of energy or
frequencies to exhibit the photoelectric
effect. (different metals are different
elements with different numbers of
electrons)
 So what did all this mean for where the
electrons were in an atom?

It was concluded that electrons exist in
specific energy levels in an atom:
 Ground state = lowest energy state of an
atom
 Excited state = the highest energy state
 When atoms are excited by energy (heat),
they emit energy in the form of light.
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Classical theory predicted that atoms
would be excited by whatever amount of
energy that was added to them. (there
would be a continuous spectrum of
frequencies given off-like a prism)
 However, when current was passed
through Hydrogen gas, a series of very
specific frequencies were emitted and only
certain colors were seen (line emission
spectrum)
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Hydrogen Line Emission
Spectrum
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This suggested that the electrons of an
atom exist in very specific energy states.
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So, Bohr put all this information together in
his model of an atom.
Bohr Model
Orbits-electrons can only circle the
nucleus in allowed paths
 Each orbit has a fixed amount of energy
 Closest to the nucleus has the least
amount of energy (ground)
 Farther from the nucleus has more energy
(excited)
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Or in other words,
When an excited electron returns to its
ground state, it gives off the energy (a
photon) in the form of electromagnetic
radiation (sometimes visible light).
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From E2 to E1, the electron will gain or
lose energy?
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From E1 to E5, the electron will gain or
lose energy?
Why do different atoms emit
different light?
Each atom is unique and contains its own
unique electron structure in the different
energy levels.
How does the emission of light
relate to the electron structure?
Since each atom is unique in its electron
structure with differing levels of energy,
the transitions between those levels will be
unique to each atom.
 Electrons are in certain energy levels.
When electrons give off light, they emit
energy, and move to a lower level closer
to the nucleus.
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Balmer Series
Emission Spectrum
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Some electron transitions result in
energies and wavelengths within the
visible light spectrum so we can see them
(400-750 nm).
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However, there are many transitions that
we cannot see (radio waves, x-rays,
gamma rays)
de Broglie concluded that since an
electron is so small but its speed is so
great, it could orbit a nucleus millions of
times in 1 second! (He used algebraic
methods and the equations of Einstein,
Planck, and the speed of a wave to figure)
 So, how could we possibly know where an
electron is in an atom?
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Schrödinger suggested that since
electrons can be thought of like waves,
they may be like standing waves outside
the nucleus.
 Only a certain number of waves can exist
between the nucleus and a certain point.
 This fits with Bohr’s idea of energy levels
in an atom.
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Ch. 4-2 The Quantum Model
Light can behave as waves and particles
 Louis De Broglie investigated that
electrons also behave like waves
because:
1. They are confined to a specific frequency
2. Diffraction-bending of a wave as it
passes by something
3. Interference-waves overlapping
Fig. 4-10
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Heisenberg Uncertainty Principle
Heisenberg: It is impossible to
simultaneously determine the location and
velocity of an electron or particle
 Schrodinger’s theory (that there can only
be so many wavelengths of energy in a
certain level) led to the development of the
quantum theory
 Quantum theory-describes the wave
patterns of electrons mathematically
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As a result of the Schrodinger equation
and Heisenberg’s Principle, the location of
an electron is only its probable location
around a nucleus
Orbital-3D region around the nucleus that
describes its probable location
 Fig. 4-11
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Quantum Numbers
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1.
2.
3.
4.
Specify the properties of atomic orbitals
and properties of electrons in orbitals
There are 4 quantum numbers for each
electron:
Principle Quantum Number
Angular Momentum Quantum Number
Magnetic Quantum Number
Spin Quantum Number
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NO two electrons have the same 4
quantum numbers
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Similar to a zip code-no 2 cities have the
same zip code
Principal Quantum Number (n)-indicates the main energy level of
the electron
Ex: n= 1, 2, 3…..
Also indicates how many sublevels there may be for a main energy
level
1.
2.
Angular Momentum Quantum Number (l)-indicates the shape of
the sublevel or oribital
S-sphere
P-dumbbell shaped
D-3D shape
F- Too complex
Pg. 102
3.
Magnetic Quantum Number (m) –
orientation of an orbital around the
nucleus
S-1
P-3
D-5
F-7
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Each orbital can hold 2 electrons
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So, total electrons for each S = 2 (1 X 2)
P = 6 (3 X 2)
D = 10 (5 X 2)
F = 14 (7 X 2)
How many electrons in each
energy level?
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Use 2n2 to figure out how many electrons
can be in each energy level
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Ex: for energy level 5, n=5
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So 2 (5)2 = 50 electrons
Spin Quantum Number – indicates the
direction the electron will spin in orbit
* has only 2 possible values for the spin,
either +1/2 or -1/2
4.
* The two electrons in each orbital have to
have opposite spins
Ch. 4-3 Electron Configurations
The quantum model tells us more than the
Bohr model of the atom because it tells us
where the electrons are located
 Electron configurations-the arrangement of
electrons in an atom
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Aufbau principle-an electron occupies the
lowest-energy orbital that can receive it
Fig. 4-16
We always start at 1s and work up to 2s, 2p,
etc.
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Pauli exclusion principle-no two electrons
in the same atom can have the same set
of four quantum numbers
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Ex: zip code
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Hund’s Rule-orbitals of equal energy are
each occupied by one electron before any
orbital is occupied by a second electron.
All the first electrons must all have the
same spin
Electron Configurations
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Electrons fill orbitals (s, p, d, f) and energy
levels (1, 2, 3…) in a certain order
according to energy:
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Remember, s orbitals can hold 2 electrons
p orbitals can hold 6 electrons
d orbitals can hold 10 electrons
f orbitals can hold 14 electrons
Electron Configurations
1. Electron Configuration Notation
Ex: He- 1s2
Ex: Na-1s22s22p63s1
2. Orbital Notation
Ex: He-_↑↓_
Ex: Na- (1s)_↑↓_ (2s)_↑↓_ (2p)_↑↓_ _↑↓_ _↑↓_
(3s) _↑_
3. Noble Gas Configuration
Ex: Li- [He] 2s1
Ex: Na-[Ne] 3s1
Practice
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Practice Configurations on pg. 114 and
116
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Or Practice writing all 3 configurations for
Al, Cu, Ag, and Cs