Transcript electron configuration

INTRODUCTORY CHEMISTRY
Concepts & Connections
Fifth Edition by Charles H. Corwin
Chapter
10
Modern Atomic
Theory &
Periodicity
Christopher G. Hamaker, Illinois State University, Normal IL
© 2008, Prentice Hall
10.1 History of the Atom
Dalton’s Model of the Atom
• John Dalton proposed that all matter is made up of
tiny particles.
• These particles are molecules or atoms.
• Molecules can be broken down into atoms by
chemical processes.
• Atoms cannot be broken down by chemical or
physical processes.
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Thomson’s Model of the Atom
• J.J. Thomson proposed a subatomic model of the
atom in 1903.
• Thomson proposed that
the electrons were
distributed evenly
throughout a homogeneous
sphere of positive charge.
• This was called the
“plum pudding” model
of the atom.
Chapter 5
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Rutherford’s Gold Foil Experiment
• Rutherford’s student fired
alpha particles at thin gold
foils. If the “plum
pudding” model was
correct, α-particles should
pass through undeflected.
• At the center of an atom is
the atomic nucleus,
which contains the atom’s
protons.
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10.2 Radiant Energy Spectrum
• The complete radiant energy spectrum is an
uninterrupted band, or continuous spectrum.
• The radiant energy spectrum includes most types of
radiation, most
of which are
invisible to
the human
eye.
Chapter 5
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Visible Spectrum
• Light usually refers to radiant energy that is visible
to the human eye.
• The visible spectrum is the range of wavelengths
between 400 and 700 nm.
• Radiant energy that has a wavelength lower than
400 nm and greater than 700 nm cannot be seen by
the human eye.
Chapter 5
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The Wave/Particle Nature of Light
• In 1900, Max Planck proposed that radiant energy
is not continuous, but is emitted in small bundles.
This is the quantum concept.
• Radiant energy has both a wave nature and a
particle nature.
• An individual
unit of light
energy is
a photon.
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Wave Nature of Light
• Light travels through space as a wave, similar to
an ocean wave.
– Wavelength is the distance light travels in one cycle.
– Frequency is the number of wave cycles completed
each second.
• Light travels at a constant speed: 3.00 × 108 m/s
(given the symbol c).
• c= ln
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Waves
Amplitude
• The distance between
corresponding points on
adjacent waves is the
wavelength (l).
• The number of waves
passing a given point per
unit of time is the
frequency (n).
Wavelength vs. Frequency
• The longer the wavelength of light, the lower the
frequency.
• The shorter the wavelength of light, the higher the
frequency.
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Sample Calculating Frequency from Wavelength
The yellow light given off by a sodium vapor lamp used for public
lighting has a wavelength of 589 nm. What is the frequency of this
radiation?
Solution
589 nm
1m
1 X 109 nm
3.00 X 108 m/s
5.89 X 10-7 m
5.89 X 10-7m
The Nature of Energy
• Einstein used this assumption to explain
the photoelectric effect. (electrons are
ejected from metals when light from
specific wavelengths are applied)
• He concluded that energy is
proportional to frequency:
E = hn
where h is Planck’s constant,
6.63  10−34 J-s.
The Nature of Energy
• Therefore, if one knows the
wavelength of light, one can
calculate the energy in one
photon, or packet, of that
light:
c = ln
E = hn
Sample Energy of a Photon
Calculate the energy of one photon of yellow light that has a
wavelength of 589 nm (n= 5.09 X 1014 s-)
Solution
The value of Planck’s constant, h, is given both in the text and in the table of
physical constants on the inside back cover of the text, and so we can easily
calculate E:
The Wave Nature of Matter
• Louis de Broglie proposed that if
light can have material
properties, matter should exhibit
wave properties.
• He demonstrated that the
relationship between mass and
wavelength was
h
l = mv
Sample Matter Waves
What is the wavelength of an electron moving with a speed of
5.97 × 106 m/s? The mass of the electron is 9.11 × 10-31 kg.
Solution
h
l = mv
mass
Kg
6.40 X 1034 J*s
(9.11 X10-31 kg)(5.97 X106 m/s)
Velocity
m/s
1.22 X 10-10 m
Emission Line Spectra
• When an electrical voltage is passed across a gas
in a sealed tube, a series of narrow lines is seen.
• These lines are the emission line spectrum. The
emission line spectrum for hydrogen gas shows
three lines: 434 nm, 486 nm, and 656 nm.
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“Atomic Fingerprints”
• The emission line spectrum of each element is
unique.
• We can use the line spectrum to identify elements
using their “atomic fingerprint.”
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10.5 Bohr Model of the Atom
• Niels Bohr speculated that electrons orbit about
the nucleus in fixed energy levels.
• Electrons are found only in specific energy levels,
and nowhere else.
• The electron energy
levels are quantized.
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The Quantum Concept
• The quantum concept states that energy is present
in small, discrete bundles.
• For example:
– A tennis ball that rolls down a ramp loses potential
energy continuously.
– A tennis ball that rolls down a staircase loses potential
energy in small bundles. The loss is quantized.
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Evidence for Energy Levels
• Bohr realized that this was the evidence he needed
to prove his theory.
• The electric charge temporarily excites an electron
to a higher orbit. When the electron drops back
down, a photon is
given off.
• The red line is the
least energetic and
corresponds to an
electron dropping
from energy level 3
to energy level 2.
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10.7 Quantum Mechanics
• Erwin Schrödinger
developed a mathematical
treatment into which both
the wave and particle
nature of matter could be
incorporated.
• This is known as quantum
mechanics.
Quantum Model
• It was later shown that electrons occupy energy
sublevels within each level.
• These sublevels are given the designations s, p, d,
and f.
– These designations are in reference to the sharp,
principal, diffuse, and fine lines in emission spectra.
• The number of sublevels in each level is the same
as the number of the main level.
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Energy Levels and Sublevels
• The first energy level has 1 sublevel:
– 1s
• The second energy level has 2 sublevels:
– 2s and 2p
• The third energy level has 3 sublevels:
– 3s, 3p, and 3d
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Electron Occupancy in Sublevels
• The maximum number of electrons in each of the
energy sublevels depends on the sublevel:
– The s sublevel holds a maximum of 2 electrons.
– The p sublevel holds a maximum of 6 electrons.
– The d sublevel holds a maximum of 10 electrons.
– The f sublevel holds a maximum of 14 electrons.
• The maximum electrons per level is obtained by
adding the maximum number of electrons in each
sublevel.
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Electrons per Energy Level
Chapter 5
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Quantum Mechanical Model
• An orbital is the region of space where there is a
high probability of finding an electron.
• In the quantum mechanical atom, orbitals are
arranged according to their size and shape.
• The higher the energy of an orbital, the larger its
size.
• s-orbitals
have a
spherical
shape
Chapter 5
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Shapes of p-Orbitals
• Recall that there are three different p sublevels.
• p-orbitals have a dumbbell shape.
• Each of the p-orbitals has the same shape, but
each is oriented along a different axis in space.
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d- orbitals
F-orbitals
10.9 Electron Configurations
• Electrons are arranged about the nucleus in a
regular manner. The first electrons fill the energy
sublevel closest to the nucleus.
• Electrons continue filling each sublevel until it is
full, and then start filling the next closest sublevel.
• A partial list of sublevels in order of increasing
energy is:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …
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Filling Diagram for Sublevels
• The order does
not strictly
follow 1, 2, 3,
etc.
• For now, use
this figure to
predict the
order of
sublevel filling.
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Orbital Diagrams
• Each box in the diagram
represents one orbital.
• Half-arrows represent the
electrons.
• The direction of the arrow
represents the relative
spin of the electron.
Electron Configurations
• The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
• The sublevel is written followed by a superscript
with the number of electrons in the sublevel.
– If the 2p sublevel contains 2 electrons, it is written 2p2.
• The electron sublevels are arranged according to
increasing energy.
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Rules for e- Configs
Electron Configurations follow 3 rules
•Aufbau- electrons fill in lowest energy first (start
at the bottom)
•Pauli Exclusion- 2 electrons maximum in an
orbital, with opposite spins to reduce repulsion
•Hund’s- everyone (box) gets one e- before
anyone gets seconds- (in degenerate orbitals)
Writing Electron Configurations
• First, determine how many electrons are in the
atom. Bromine has 35 electrons.
• Arrange the energy sublevels according to
increasing energy:
– 1s 2s 2p 3s 3p 4s 3d …
• Fill each sublevel with electrons until you have
used all the electrons in the atom:
– Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p5
• The sum of the superscripts equals the atomic
number of bromine (35).
Chapter 5
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Valence Electrons
• When an atom undergoes a chemical reaction,
only the outermost electrons are involved.
• These electrons are of the highest energy and are
furthest away from the nucleus. These are the
valence electrons.
• The valence electrons are the s and p electrons
beyond the noble gas core.
• For our purposes we will include ALL valence
electrons past the noble gas core.
10.10 e- Configs using the Periodic Table
• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table (shaded in
different colors in this chart) correspond to
different types of orbitals.
Blocks and Sublevels
• We can use the periodic table to predict which
sublevel is being filled by a particular element.
Noble Gas Core Electron Configuration
• Recall, the electron configuration for Na is:
Na: 1s2 2s2 2p6 3s1
• We can abbreviate the e- config by indicating the
innermost electrons with the symbol of the
preceding noble gas.
• The preceding noble gas before sodium is neon,
Ne. We rewrite the electron configuration:
Na: [Ne] 3s1