Transcript Chapter Two Atoms & The Periodic Table

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Ability to do work or transfer heat
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Kinetic-Energy due to motion
 Thermal
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Potential-Energy due to position/arrangement
 Chemical
 Electrostatic
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Energy is measured in Joules (J); N*m; kg*m2/s2
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Last time in Chemistry…
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Atoms had protons, neutrons, & electrons
 Protons and neutrons in dense core called nucleus
 Electrons orbit nucleus (provide most of volume)
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The atom was assumed to look a little bit like
a microscopic solar system
 Classical descriptions:
◦ Dalton: atoms are hard particles, all atoms of the same
element are the same
◦ Energy is continuous
◦ Planetary model of atom
◦ Newtonian laws (macroscopic) were not working for
atomic behavior (microscopic)
◦ Scientists used interactions between matter and light to
gain insight on atomic structure
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Light is typically described as traveling in waves (similar to water)
All electromagnetic (EM) waves are made of two components:
electric waves and magnetic waves
ALL EM waves travel at the speed of light, c (2.998 x 108 m/s)
 c = ln
c = speed of light; l (lambda)= wavelength (m, nm);
ν (nu) = frequency (1/s, s-1, Hz)
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Two wavelengths (l) are shown below.
Determine the frequency (n) for each wave.
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1 nm = 1 x 10-9 m
OR 1 x 109 nm = 1 m
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Determine what color each wave below
would be. Then determine the frequency (v)
of wave (a).
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Light exhibits wave AND particle properties
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Double Slit Experiment
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When light passes through two closely spaced
slits, an interference pattern is produced.
This type of
interference is
typical of waves
and demonstrates
the wave nature of
light
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Light makes impressions on screen, suggesting
light has particle-like behavior
 Blackbody Radiation: The visible glow solid objects
give off when heated
▪
Intensity varies with wavelength, λ, of emitted light
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Classical physics failed to explain this
phenomenon
 Assumed radiant energy was continuous (could be
emitted or absorbed in any amount)
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Max Planck suggested radiant energy is only
emitted or absorbed in discrete quantities
called quanta
E = hν
Planck’s constant, h = 6.626 x 10-34 J·s
Energy, E, in SI unit of joule (kg·m2/s2)
The Photoelectric Effect—Albert Einstein:
used Planck’s theory to observe metals
reacting to different colors of light –electrons
are ejected from the surface of certain metals
exposed to light at a certain minimum
frequency
 Energy of photon
based solely on
frequency NOT intensity
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1) Which has a higher frequency: light from a red stoplight
with a wavelength of 750 nm or a yellow light with a
wavelength of 600 nm?
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2) What is the wavelength of a radio station’s waves
transmitting at a frequency of 101.5 MHz (megahertz)?
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3) Calculate the energy (in Joules) of a photon of light with a
wavelength of 5.00 x 104 nm (infrared region).
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Answers: yellow, 2.956 m, 3.98 x 10-21 J
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Quantum has come to mean small;
originated from Planck’s observation of
quantized energy
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New set of “rules” had to be used for the
subatomic world
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As electrons orbit, they radiate energy (due
to acceleration)
If electrons radiate energy, they must
eventually start spiraling towards nucleus
Electron would eventually hit the nucleus
where the proton is and the atom would
annihilate itself
This doesn’t happen…whew!
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Neils Bohr proposed the idea that electrons
orbit the nucleus in FIXED orbits
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Each orbit is a different energy level
Electrons in different orbits have different
energies
Electrons cannot exist between energy levels
n = 1 is the lowest energy level
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Electrons travel in discrete, quantized circular
orbits; like going up or down stairs.
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If an electron is to “jump” orbits (move away
from the nucleus), it needs to ABSORB
energy in just the right quantity
If an electron is to “drop” orbits (move closer
to the nucleus), it will RELEASE energy in the
form of electromagnetic radiation
Sometimes visible light, sometimes not
The Hydrogen Atom
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Hydrogen 
Each element gives
off unique spectrum
Spectra of Elements:
http://www.wwnorton.com/college/chemistry/chemconnections/BlueLight/pages/elements.html
Figure 7.8
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Bohr stated that electrons move in FIXED
orbits.
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But why???
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de Broglie: If light (traditionally believed to be a
wave) can behave like a particle, then electrons
(traditionally believed to be matter) can behave
like a wave
If an electron behaves like a standing wave, then
it can only have specific wavelengths
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Can calculate wavelength for ANY matter if
we know its velocity (use v instead of c):
l = h / m v (This is the de Broglie equation.)
 h = Planck’s constant, m = mass (electron’s have
constant mass: 9.11x10-31 kg), v = velocity (speed
in m/s)
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Based on photoelectric effect, light acts as a
wave but also exists as a stream of particles
called photons
Energy of photons is proportional to
frequency, inversely proportional to
wavelength
E  hn 
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hc
l
J = kg • m2 / s2
h = 6.626 x 10-34 J•s
1) What is the wavelength of an electron that travels at
34.7 m/s and has a mass of 9.11 x 10-31 kg?
2) A 0.143 kg baseball is thrown at a velocity of 42.5 m/s.
Calculate the wavelength of the baseball.
3) How do the wavelengths compare from 1 & 2?
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If electrons have wave-like properties and
particle-like, then we can’t know both its
position and velocity (momentum) at the
same time
 In order to determine the position of an electron,
we hit it with a photon of light, but this will
change its position and velocity.
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Quantum Mechanics was developed (by
Schrödinger in the 1920’s) to describe the
motion of subatomic particles
 Did not attempt to describe exact position of
particles; used mathematical equations to
describe the probability of finding the particles
 The probability density (map of likely locations)
is the “electron cloud”
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The region of highest probability for finding
an electron is an “electron cloud”. This region
of high probability is also called an atomic
orbital. Each orbital holds at most 2
electrons.
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There are 4 quantum number that help
describe the orbitals for electrons
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We use these numbers to describe where
electrons are most likely to be found for an
atom.
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n: shells (distance from nucleus)
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l: subshells (shape)
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ml: orbitals (orientation)
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ms: spin (electron spin)
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The Principal Quantum Number, n
 describes distance of the electron from the
nucleus; called shells
 n = 1, 2, 3, etc; larger number is farther from
nucleus
 n typically corresponds to row # in the periodic
table
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The Angular Momentum Quantum Number, l
 Refers to the shape of the orbitals.
 These groups of orbitals are called subshells and
labeled s (0), p (1), d (2), and f (3).
▪ l = n-1
 s subshells are spherical (first two columns)
 p subshells are dumb-bell shaped (last six
columns)
 d subshells are intersecting dumb-bells
(transition metals)
 f subshells are lanthanides/actinides
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Arrangement of subshells in the Periodic
Table
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The Magnetic Quantum Number, ml
 describes the orientation of the orbital with
respect to x, y, and z axes
 -l …. +l
The Spin Quantum Number, ms
 describes the spin of an electron in an
orbital (shown as up and down arrows in
orbital diagrams)
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Orbitals of Scandium
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What is the maximum number of:
 electrons allowed in the 2px orbital?
 subshells allowed in the 4th shell?
 electrons allowed in the 3d subshell?
 electrons allowed in the 4d subshell?
 electrons allowed in the 3p subshell?
 electrons allowed in the 3rd shell?
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In hydrogen, all
shells are equivalent
in energy.
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In many-electron
models, the energy
levels depend on the
shell and subshell.
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Aufbau principle:
electrons fill from
lowest energy level to
highest energy level
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Write electron
configurations for the
following atoms.
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Al
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Fe
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Br
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U
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Rather than writing out complete electron
configurations, we can use the previously
filled shell (noble gas) and show the valence
electrons (v. e.):
P: 1s2 2s2 2p6 3s2 3p3  [Ne] 3s2 3p3 (5 v. e.)
Write the shorthand notation for:
 Pd
 Pu
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Some exceptions to the Aufbau order…
What are the expected electron
configurations for Cr and Cu?
Filled and half-filled d subshells seem to be
especially stable.
Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
 Also true for Mo and W
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Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
 Also true for Ag and Au
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Electrons in the outermost shell.
 1s2 2s2 2p6
 1s2 2s2 2p6 3s2 3p5
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Identify the number of valence electrons (v.
e.) in the following configurations:
 1s2 2s2 2p6 3s2
 1s2 2s2 2p63s2 3p6 4s2 3d10 4p4
1s2 2s2
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Write electron configurations for the
following elements (long-hand notation).
 Indicate the number of v.e. for each
element.
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Potassium
Gold
Silver
Neon
Orbital
diagrams are
pictorial
representations
of electron
configurations.
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Selenium (Se)
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If two or more orbitals (i.e., a p orbital) with the
same energy are available, one electron goes into
each orbital until they have to pair up.
 Fighting siblings
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For example, an atom with 4 p electrons: 1 electron
will go into the first (px) orbital, the next electron
will go into the second (py) orbital, the 3rd will go in
the (Pz ) orbital, and THEN the 4th will fill the (Px)
orbital, making it full
Also, all electrons in singly-occupied orbitals must
have same spin
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Pauli Exclusion Principle: no two electrons
can have the same values of all 4 quantum
numbers
Describes what happens when electrons
share an orbital.
 Only two electrons can occupy a single orbital and
they must have opposite spin (i.e., the 4th
quantum number). The first electron is
designated as positive spin (up arrow), the second
electron in that orbital has negative spin (down
arrow).
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Portions of orbital diagrams are shown below for
atoms in their ground state configurations. For each
letter, state whether it violates Hund’s Rule or Pauli
Exclusion Principle (one of the violates neither!)