Chemistry I Electrons in Atoms Chapter 5

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Chemistry IElectrons in AtomsChapter 5
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Rutherfords nuclear model did not provide enough
detail about how electrons occupy the space
around the nucleus. In this chapter we will learn
how electrons are arranged around the nucleus and
how that arrangement effects chemical behavior.
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In what ways did many scientists
findRutherfords model to be incomplete?
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1.)In what ways did many scientists
findRutherfords model to be incomplete?

• It did not explain
• 1.How the electrons were arranged around the
  nucleus
• 2.Why they were not pulled into the nucleus
• 3. The differences in chemical behavior of the
  different elements.

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2.)In the early 1900s what did scientists
observe about certain elements when heated in a
flame?

• Every element gives off specific wavelengths of
  light that can be used like a fingerprint to
  identify the element.

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spectroscope
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• http//jersey.uoregon.edu/vlab/elements/Elements.h
  tml

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An elements chemical behavior is related to the
arrangement of electrons in its atoms.
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Wave Nature of Light
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What is electromagnetic radiation?
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3.)What is electromagnetic radiation?

• A type of wave that is part electrical and
  magnetic energy acting at right angles.
• Visible light is a type of electromagnetic
  radiation.

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What are the 4 characteristics of waves?
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4.) What are the 4 characteristics of waves?

• 1.) wavelength distance from crest to crest
• 2.) frequency- of waves that pass a given point
  per second
• 3.) Amplitude height of wave from the normal
  resting to wave crest
• 4.) speed the speed of light is a constant 3.0
  x 108 m/s.

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5.)What unit do we use to express frequency?

• Hertz (Hz)
• The hertz unit is 1/second ( the inverse of a
  second)

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6.) What is the speed of light?

• The speed of light is a constant 3.0 x 108 m/s
• The formula for the speed of light is C ? v
• ? wavelength (measured in meters)
• v frequency (measured in 1/sec)
• The unit for speed is m/s (distance/time)

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6 continued

• When wavelength (?) increases then frequency (v)
  decreases because the speed of light ( c) is a
  constant 3.0X108m/s.
• Also
• When ? decreases, V increases

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7.) List types of electromagnetic radiation from
the lowest to the highest in energy ( page 139)
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List types of electromagnetic radiation from the
lowest to the highest in energy ( page 140)

• Radio waves ?decreasing wavelength
• Microwaves ?increasing frequency
• Infrared
• Visible
• Ultraviolet (U.V.)
• X rays
• gamma

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Particle Nature of Light

• When an object is heated only certain
  wavelengths of light are emitted. The instrument
  used to separate light into its different
  wavelengths is called a spectroscope. The pattern
  of wavelengths is called the spectrum.

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• http//jersey.uoregon.edu/vlab/elements/Elements.h
  tml

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• When a substance is heated the spectrum of light
  that is given off can be used to identify the
  substance.

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• http//jersey.uoregon.edu/vlab/elements/Elements.h
  tml

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• The wave model of light could not explain why
  different substances emit particular wavelengths
  of light when heated

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8.)What are some problems with the wave model of
the atom?

• It doesnt explain why only certain wavelengths
  of light are emitted when an object is heated.
• It also does not explain the photoelectric
  effect.( we will discuss the photoelectric effect
  in just a little bit)

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Who is Max Planck?
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9.)Who is Max Planck?

• A German physicist who, in 1900, began searching
  for the reason that only certain wavelengths of
  light are given off for every element.

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10.)What did Planck conclude?

• Matter can gain or lose energy in only small
  specific amounts called quanta.

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As metal is heated, it glows. In fact, it glows
different colors as it gets hotter (white hot
being the hottest). By studying glowing metals,
Max Planck discovered that only certain
wavelengths of light are emitted at each specific
temperature.
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11.) What is a quantum?

• The minimum amount of energy that can be gained
  or lost by an atom.
• The amount of energy it takes for a electron to
  get from one energy level to the next.

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12.) If energy can only absorbed in quanta,
specific amounts, why does energy appear to
continuous to us?

• b/c quanta are extremely small.

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13.) What is Plancks equation that demonstrates
that energy is related to the frequency of
radiation?

• Equation
• E hv
• E energy of a quantum
• h- plancks constant
• v- frequency

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14.) What is Plancks constant?

• 6.62 x 10-34 Js

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15.) How are energy and frequency related?

• They are directly related
• As frequency increase, energy increases.

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What is the photoelectric effect?
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What is the photoelectric effect?
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• http//www.tutorvista.com/content/physics/physics-
  iv/radiation-and-matter/photoelectric-effect-and-c
  ell.php

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16.)What is the photoelectric effect?

• When light of a certain frequency shines on the
  surface of a metal, electrons are ejected. Blue
  light always results in electrons being ejected.
  Red light of any brightness does not eject
  electrons.

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17.) In 1905 how did Albert Einstein explain the
photoelectric effect?

• He proposed that all electromagnetic radiation
  (including visible light) is both wavelike and
  particle like in nature.
• Continued next slide

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17.) In 1905 how did Albert Einstein explain the
photoelectric effect?

• He suggested that while a beam of light has many
  wavelike characteristics, it is also like a
  stream of particles called photons.

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18.) What is a photon?

• A photon is a particle of electromagnetic
  radiation with no mass that carries a quantum of
  energy.

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19.)How did Einstein use Plancks idea of quantum
energy to explain the photoelectric effect?

• Planck proposed that Ehv. Energy is equal to his
  constant times the frequency of radiation.
• Einstein proposed that since blue light has a
  larger frequency, it has a larger quantum of
  energy.
• Continued next slide

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19.) How did Einstein use Plancks idea of
quantum energy to explain the photoelectric
effect?

• E hv for blue light is more than Ehv for red
  light because blue light has a larger frequency
  than red light.

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Practice Problems page 143
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20.)What is the atomic emission spectrum?

• The atomic emission spectrum of an element is the
  set of frequencies of (light) electromagnetic
  waves emitted ( given off) by atoms to that
  element when it is heated.

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21.) How does the atomic emission spectrum for
one element compare to another element?

• Each is unique and can be used to identify the
  element like a fingerprint.

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22.) Are atomic emission spectrums of elements
continuous or distinct individual wavelengths of
light?

• They are distinct, individual wavelengths of
  light.

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23.) In a line spectrum which line has the
highest energy?

• The line farthest to the blue end of the
  spectrum. The line corresponding to the shortest
  wavelength/ highest frequency.

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24.) How does the quantitization of energy (
energy only exists in specific sizes) help
explain line spectrum?

• The energy of the photon of light that is emitted
  is tied to a specific frequency/color of light
• E hv Each frequency (v) corresponds to a
  specific color

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Questions page 145
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Section 2 Quantum Theory and the Atom
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• Scientists concluded that light was both wave and
  particle. Scientists were able to understand
  atomic structure, electrons and atomic emission
  spectra better because of a better understanding
  of light.

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• Scientists wanted to know why atomic emission
  spectrums are not continuous but are
  discontinuous.

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23.)Who was Niels Bohr?

• A Danish physicist who worked in Rutherfords
  laboratory. He proposed the quantum model of the
  atom that helped explain why atomic emission
  spectrums only contain certain frequencies of
  light.

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25.)What did Niels Bohr propose?

• He proposed that a hydrogen atom has only certain
  allowable energy levels.
• The lowest level is called the ground state.
• continued

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25.)What did Niels Bohr propose?

• All other levels are called excited states.
• Electrons move around the nucleus in only certain
  allowed orbits.
• Bohr labeled these orbits.
• continued

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25.)What did Niels Bohr propose?

• The orbit closest to the nucleus is labeled n1
  and was lowest in energy.

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26.) How did Niels Bohr explain line spectrum
with the quantization of energy for electrons. (
How did Bohr explain that certain wavelengths of
light are released when an element is heated by
proposing that electrons can only be at certain
energy levels.)
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• He said that when an element is at ground state (
  the lowest energy level) no energy (light) is
  released. But when an electron becomes excited,
  it jumps up to another energy level.

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• When the electron returns to the lower energy
  level the energy is released. The energy (light)
  that is released is equal to the difference
  between energy levels.

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• The amount of energy that is released is equal to
  a particular frequency of light.

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27.)Do we still accept Bohrs model?

• Bohr was correct about his idea of quantization
  of energy and his basic explanation of atomic
  emission spectrum. However he could only predict
  the atomic emission spectrum of hydrogen.

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27.)Do we still accept Bohrs model?

• His basic model of the model is considered to be
  incorrect. Today we do not believe that the
  electrons are in orbits so that we can predict
  their positions.

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The Quantum Mechanical Model of the Atom
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28.)Who was Louis de Broglie?

• A French physics student that proposed an idea
  that eventually accounted for the fixed energy
  levels in Bohrs model.

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29.)What did De Broglie propose?

• He proposed that all matter moves in waves. The
  larger the mass the smaller the wavelength. Only
  very small objects have noticeable wavelengths.

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30.)What was De Broglies equation?

• ? h/ mv
• DeBroglies equation predicts a large mass will
  have a small wavelength (?) only a very small
  mass will have a noticeable wavelength.

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Page 149
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If electrons move in waves then only electrons
with matching wavelengths can be accepted into an
energy level and when an electron is released to
a lower energy level particular wavelengths of
energy are released.
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29.)What did Heisenberg propose?

• It is impossible to make a measurement of an
  object without disturbing it. For example, a
  thermometer changes the temperature of an object.

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30.)What is the Heisenberg uncertainty principle?

• It is impossible to know precisely both the
  velocity and position of a particle at the same
  time.

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The Schrodinger Wave Equation
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32.) The Austrian physicist Edwin Schrodinger
derived an equation that treated the electron as
a wave. What was significant about his equation?

• It worked equally well for atoms of other
  elements unlike Bohrs model which worked only
  for hydrogen.

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33.) What is the name of the model proposed in
which electrons are treated as waves?

• The quantum mechanical model of the atom.

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34.)What is the quantum mechanical model of the
atom?

• It limits the electrons energy to certain values
  but makes no attempt to describe the path of the
  electron around the nucleus

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35.) What is an atomic orbital?

• In the quantum mechanical model of the atom, the
  orbital is the region in which we expect to find
  an electron 90 of the time.
• Orbitals are described by electron clouds.

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Hydrogens Atomic Orbital
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36.)What are principal quantum numbers?

• Principal quantum numbers represent the relative
  sizes and energies of the atomic orbitals.
• As n increases, the size and energy level
  increases.

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37.) What are sublevels?

• Each principle energy level has sublevels. The
  number of sublevels is equal to the n level.
• n 1 has one sublevel
• n 2 has two sublevels
• etc.

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38.) What type of sublevels are there?

• The sublevels are named s, p, d, and f

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Sublevel Shape energy level

• s spherical lowest
• p dumbell
• d cloverleaf
• f complex highest

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39.)How many electrons does an orbital contain?

• two

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40.) How many orbitals does each sublevel
contains?

• Sublevel orbitals e-
• s 1 2
• p 3 6
• d 5 10
• f 7 14

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n 4n 3n 2n 1
s p d f
s p d
s p
s
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Principal Quantum Sublevel types orbitals Total of orbitals Total electrons
1 s 1 1 2
2 s p 1 3 4 8
3 s P d 1 3 5 9 18
4-7 s p d f 1 3 5 7 16 32

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Section 3 Electron ConfigurationsThe arrangement
of electrons in atoms follows a few very specific
rules.
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Ground State Electron Configurations
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41.)What is an electron configuration?

• The arrangement of electrons in an atom

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42.) What is ground state electron configuration?

• The most stable, lowest energy arrangement of
  electrons.

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43.) What are the three rules that govern ground
state electron configurations.

• 1.) Aufbau principle
• 2.) Hunds Rule
• 3.) Pauli Exclusion Principle

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44.)What is the Aufbau principle?

• Electrons occupy the lowest energy orbital
  available.
• German for building up

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45.) In general what is the order of increasing
energy among energy levels and sublevels?

• In general n1 is lower than n2 . The order of
  increasing order of sublevels is s, p, d, f.
• Use the diagram

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46.) What is the Pauli Exclusion Principle?

• A maximum of two electrons can occupy an orbital
  and they must have opposite spins.
• The way to indicate two electrons with opposite
  spin is

??
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47.) What is the Hunds Rule?

• Single electrons with the same spin must occupy
  equal energy orbitals before additional electrons
  with opposite spin can occupy the same orbital.

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48.)What are the two ways that electron
configurations can be described?

• Orbital diagrams using boxes
• Mg
• 1s 2s 2p 3s

??
??
??
??
??
??
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49.) What is noble gas configuration?

• Electron configurations that are used to show
  just the valence electrons. The full inner core
  orbitals are represented by the noble gas symbol
  with the lower atomic number and the electrons in
  the valence shell.
• K 1s22s22p63s23p64s1
• K Ar 4s1

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50.) Exceptions to predicted configurations.

• Cr Ar4s13d5
• Cu Ar4s13d10
• It is more stable to borrow an electron from
  the s orbital and put it in the d orbital

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Pg 160 practice problems
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51.) Which electrons determine the chemical
properties of an element?

• The valence electrons/ the outermost electrons.

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52. How does the number of valence electrons
compare to the family the element is in?

• Group 1 1 valence electron
• Group 2 2 valence electrons
• Group 13 3 valence electrons
• Group 14 4 valence electrons
• Group 15 5 valence electrons
• Group 16 - 6 valence electrons
• Group 17 7 valence electrons
• Group 18 8 valence electrons except for
  hydrogen which only has two.