Ch. 5: Electrons in Atoms

1
Ch. 5 Electrons in Atoms

• 5.1 Models of Atomstimeline, Bohr, Schrodinger,
  atomic orbitals
• 5.2 Electron Arrangement in Atomselectron
  configurations
• 5.3 Physics and the Quantum Mechanical
  Modellight, electromagnetic radiation, atomic
  spectra
• Vocab quantum, atomic emission spectrum, ground
  state, photon

2
Review Properties of Electric Charge

• Electric charge may be only () or (-).
• Opposite charges attract each other

Like charges repel.
3
Review Democritus (400 B.C.) and John Dalton,
1803

• Dalton
• Each element consists of a particular kind of
  atom. All atoms of a particular element are
  identical.
• His atoms had a definite sizethey were just
  really small.

Democritus Everything is made up of a few
simple parts called atomos. Atomos means
uncuttable in Greek.He envisioned atomos as
small, solid particles of many different sizes
and shapes.
4
Review Plum Pudding (Thomson) and Rutherford
Model

• Atoms contain negative particles called
  electrons and the rest of the atom is positively
  charged.

The mass of the atom is contained in a tiny
positive nucleus, which is surrounded by
electrons moving at high speeds.
5
5.1 Models of the Atom
5.1

• The scale model shown is a physical model.
  However, not all models are physical. In fact,
  several theoretical models of the atom have been
  developed over the last few hundred years. You
  will learn about the currently accepted model of
  how electrons behave in atoms.

6
The Development of Atomic Models
5.1
This timeline needs to be in your notes! 5.1
Models of the Atom

• The timeline shoes the development of atomic
  models from 1803 to 1911.

DALTON----THOMSON (PLUM PUDDING)RUTHERFORD
(NUCLEUS)
7
The Development of Atomic Models
5.1

• The timeline shows the development of atomic
  models from 1913 to 1932.

BOHR (ELECTRONS ORBIT)DE BROGLIE (WAVE
MOVEMENT)SCHRODINGER (ELECTRON CLOUD)
8
The Development of Atomic Models
5.1

• What was inadequate about Rutherfords atomic
  model?
• Rutherfords atomic model could not explain the
  chemical properties of elements.
• Rutherfords atomic model could not explain why
  objects change color when heated.
• Why objects when heated
• to higher and higher temp.first glow red,
  thenyellow, then white

9
Niels Bohr, 1913 (1911?)

• The electrons circle the nucleus but are
  restricted to particular orbits, like the planets
  around the Sun.
• Proved this using calculations of different
  energies of the atom.

1) An electron is found only in specific
circular paths, or orbits, around the nucleus 2)
Each orbit has a fixed energy called energy
levels 3) A quantum of energy is the amount of
energy required to move an electron from one
energy level to another
10
The Bohr Model
5.1

• What was the new proposal in the Bohr model of
  the atom?
• Bohr proposed that an electron is found only in
  specific circular paths, or orbits, around the
  nucleus.
• Each possible electron orbit in Bohrs model has
  a fixed energy.
• The fixed energies an electron can have are
  called energy levels.
• A quantum of energy is the amount of energy
  required to move an electron from one energy
  level to another energy level.

11
The Bohr Model

• Bohr proposed that an electron is found only in
  specific circular paths, or orbits, around the
  nucleus
• Each orbit has a particular energy level
  electrons must gain or lose energy to move from
  one energy level to the next
• A quantum of energy is the amount of energy
  required to move an electron from one energy
  level to another energy level

12
The Bohr Model
5.1

• Like the rungs of the strange ladder, the energy
  levels in an atom are not equally spaced.
• The higher the energy level occupied by an
  electron, the less energy it takes to move from
  that energy level to the next higher energy level.

13
Erwin Schrödinger, 1926 (1922?)

• The electrons do not orbit the nucleus.
  Instead, their position is determined by
  probability. 95 of the time, they can be found
  in Bohrs proposed orbits. But not always!
• Proved this using calculations.

The quantum mechanical model determines the
allowed energies an electron can have and how
likely it is to find the electron in various
locations around the nucleus (electrons are found
in orbitals drawn as an electron cloud)
14
The Quantum Mechanical Model
5.1

• What does the quantum mechanical model determine
  about the electrons in an atom?
• The quantum mechanical model determines the
  allowed energies an electron can have and how
  likely it is to find the electron in various
  locations around the nucleus.
• Austrian physicist Erwin Schrödinger (18871961)
  used new theoretical calculations and results to
  devise and solve a mathematical equation
  describing the behavior of the electron in a
  hydrogen atom.
• The modern description of the electrons in atoms,
  the quantum mechanical model, comes from the
  mathematical solutions to the Schrödinger
  equation.

15
the Quantum Mechanical Model

• Mathematically determined by Erwin Schrodinger
• Determines the allowed energies an electron can
  have and how likely it is to find the electron in
  various locations around the nucleus
• An atomic orbital is often thought of as a region
  of space in which there is a high probability of
  finding an electron.
• Each energy sublevel corresponds to an orbital of
  a different shape, which describes where the
  electron is likely to be found.
• Different atomic orbitals are denoted by letters.
  The s orbitals are spherical, and p orbitals are
  dumbbell-shaped. (s, p, d, and f )

16
Atomic Orbitals
5.1

• Four of the five d orbitals have the same shape
  but different orientations in space.

The numbers and kinds of atomic orbitals depend
on the energy sublevel.
17
The Quantum Mechanical Model
5.1

• The propeller blade has the same probability of
  being anywhere in the blurry region, but you
  cannot tell its location at any instant. The
  electron cloud of an atom can be compared to a
  spinning airplane propeller.

In the quantum mechanical model, the probability
of finding an electron within a certain volume of
space surrounding the nucleus can be represented
as a fuzzy cloud. The cloud is more dense where
the probability of finding the electron is high.
18
Atomic Orbitals
5.1

• The number of electrons allowed in each of the
  first four energy levels are shown here.

19
Using the periodic table
6.2

20
5.1 Section Quiz.

• 1. Rutherford's planetary model of the atom could
  not explain
• a) any properties of elements.
• b) the chemical properties of elements.
• c) the distribution of mass in an atom.
• d) the distribution of positive and negative
  charges in an atom.

21
5.1 Section Quiz.

• 2. Bohr's model of the atom proposed that
  electrons are found
• a) embedded in a sphere of positive charge.
• b) in fixed positions surrounding the nucleus.
• c) in circular orbits at fixed distances from the
  nucleus.
• d) orbiting the nucleus in a single fixed
  circular path.

22
5.1 Section Quiz.

• 3. What is the lowest-numbered principal energy
  level in which p orbitals are found?
• a) 1
• b) 2
• c) 3
• d) 4

23
5.2 Electron Arrangement in Atoms
5.2

• If this rock were to tumble over, it would end
  up at a lower height. It would have less energy
  than before, but its position would be more
  stable. You will learn that energy and stability
  play an important role in determining how
  electrons are configured in an atom.

24
Electron Configurations
5.2

• The ways in which electrons are arranged in
  various orbitals around the nuclei of atoms are
  called electron configurations.Three rulesthe
  aufbau principle, the Pauli exclusion principle,
  and Hunds ruletell you how to find the electron
  configurations of atoms.

• You will need to know electron configurations of
  elements with atomic number 1-20
• ex calcium is 1s2 2s2 2p6 3s2 3p6 4s2

25
Electron Configurations
5.2

• Aufbau Principle
• According to the aufbau principle, electrons
  occupy the orbitals of lowest energy first. In
  the aufbau diagram below, each box represents an
  atomic orbital.

1s 2s, 2p 3s, 3p 4s, 3d, 4p 5s, 4d, 5p 6s, 4f,
5d, 6p 7s, 5f, 6d, 7p
26
Electron Configurations
5.2

• Pauli Exclusion Principle
• According to the Pauli exclusion principle, an
  atomic orbital may describe at most two
  electrons. To occupy the same orbital, two
  electrons must have opposite spins that is, the
  electron spins must be paired.
• Hunds Rule
• Hunds rule states that electrons occupy
  orbitals of the same energy in a way that makes
  the number of electrons with the same spin
  direction as large as possible.

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Electron Configurations
5.2
According to the aufbau principle, electrons
occupy the orbitals of lowest energy first.
According to the Pauli exclusion principle, an
atomic orbital may describe at most two
electrons. To occupy the same orbital, two
electrons must have opposite spins
Hunds rule states that electrons occupy orbitals
of the same energy in a way that makes the number
of electrons with the same spin direction as
large as possible.
28
Electron Configuration and Orbital Filling
Practice

• Write the electron configuration and orbital
  filling for
• a. Li
• b. Mg
• c. Si

• Li atomic number 3
• 1s22s1

b. Mg atomic number 12
1s22s22p63s2
1s22s22p63s23p2

• Si atomic number 14

29
Electron Configurations in Groups

Group 1A elements -- there is only one electron
in the highest occupied energy level.
Group 4A elements -- there are four electrons in
the highest occupied energy level.
noble gases are the elements in Group 8A
30
Exceptional Electron Configurations
5.2

• Why do actual electron configurations for some
  elements differ from those assigned using the
  aufbau principle?
• Some actual electron configurations differ from
  those assigned using the aufbau principle because
  half-filled sublevels are not as stable as filled
  sublevels, but they are more stable than other
  configurations.
• Exceptions to the aufbau principle are due to
  subtle electron-electron interactions in orbitals
  with very similar energies.
• Copper has an electron configuration that is an
  exception to the aufbau principle.
• It is 1s22s22p63s23p64s13d10
• Instead of 1s22s22p63s23p64s23d9

31
5.2 Section Quiz.

• 1. Identify the element that corresponds to the
  following electron configuration 1s22s22p5.
• a) F
• b) Cl
• c) Ne
• d) O

32
5.2 Section Quiz.

• 2. Write the electron configuration for the atom
  N.
• a) 1s22s22p5
• b) 1s22s22p3
• c) 1s22s1p2
• d) 1s22s22p1

33
5.2 Section Quiz.

• 3. The electron configurations for some elements
  differ from those predicted by the aufbau
  principle because the
• a) the lowest energy level is completely filled.
• b) none of the energy levels are completely
  filled.
• c) half-filled sublevels are less stable than
  filled energy levels.
• d) half-filled sublevels are more stable than
  some other arrangements.

34
5.3 Physics and the Quantum Mechanical Model
5.3

• Neon advertising signs are formed from glass
  tubes bent in various shapes. An electric current
  passing through the gas in each glass tube makes
  the gas glow with its own characteristic color.
  You will learn why each gas glows with a specific
  color of light.

35
Light
5.3

• How are the wavelength and frequency of light
  related?
• The amplitude of a wave is the waves height from
  zero to the crest.
• The wavelength, represented by ? (the Greek
  letter lambda), is the distance between the
  crests.
• The frequency, represented by ? (the Greek letter
  nu), is the number of wave cycles to pass a given
  point per unit of time.
• The SI unit of cycles per second is called a
  hertz (Hz).
• The wavelength and frequency of light are
  inversely proportional to each other.

36
Light
5.3

• The product of the frequency and wavelength
  always equals a constant (c), the speed of light.

37
Light
5.3

• According to the wave model, light consists of
  electromagnetic waves.
• Electromagnetic radiation includes radio waves,
  microwaves, infrared waves, visible light,
  ultraviolet waves, X-rays, and gamma rays.
• All electromagnetic waves travel in a vacuum at a
  speed of 2.998 ? 108 m/s.
• Sunlight consists of light with a continuous
  range of wavelengths and frequencies.
• When sunlight passes through a prism, the
  different frequencies separate into a spectrum of
  colors.
• In the visible spectrum, red light has the
  longest wavelength and the lowest frequency.

38
Light
5.3

• Sunlight consists of light with a continuous
  range of wavelengths and frequencies.
• When sunlight passes through a prism, the
  different frequencies separate into a spectrum of
  colors.
• In the visible spectrum, red light has the
  longest wavelength and the lowest frequency.
• ROYGBIV (red, orange, yellow, green, blue,
  indigo, violet)

39
Atomic Spectra
5.3

• A prism separates light into the colors it
  contains. When white light passes through a
  prism, it produces a rainbow of colors.

When atoms absorb energy, electrons move into
higher energy levels. The atom becomes unstable
so the electrons then lose energy by emitting
light, or photons, when they return to lower
energy levels.
40
Atomic Spectra

• The frequencies of light emitted by an element
  separate into discrete lines to give the atomic
  emission spectrum of the element.

When light from a helium lamp passes through a
prism, discrete lines are produced.
Mercury
Nitrogen
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
electron.
41
An Explanation of Atomic Spectra

• In the Bohr model, the lone electron in the
  hydrogen atom can have only certain specific
  energies.
• When the electron has its lowest possible energy,
  the atom is in its ground state.
• Excitation of the electron by absorbing energy
  raises the atom from the ground state to an
  excited state.
• A quantum of energy in the form of light (photon)
  is emitted when the electron drops back to a
  lower energy level.

42
Quantum Mechanics
5.3

• How does quantum mechanics differ from classical
  mechanics?
• CLASSICAL In 1905, Albert Einstein successfully
  explained experimental data by proposing that
  light could be described as quanta of energy.
• The quanta behave as if they were particles.
• Light quanta are called photons.
• Ex. Electron absorbing energy from ground state
  and then releasing quanta of energy to return to
  ground state
• QUANTUM In 1924, De Broglie developed an
  equation that predicts that all moving objects
  have wavelike behavior.

43
Quantum Mechanics
5.3

• Today, the wavelike properties of beams of
  electrons are useful in magnifying objects. The
  electrons in an electron microscope have much
  smaller wavelengths than visible light. This
  allows a much clearer enlarged image of a very
  small object, such as this mite.

Classical mechanics adequately describes the
motions of bodies much larger than atoms, while
quantum mechanics describes the motions of
subatomic particles and atoms as waves.
44
Quantum Mechanics
5.3

• The Heisenberg uncertainty principle states that
  it is impossible to know exactly both the
  velocity and the position of a particle at the
  same time. Once youve located the particle, the
  particle has moved.
• This limitation is critical in dealing with small
  particles such as electrons.
• This limitation does not matter for
  ordinary-sized object such as cars or airplanes.
• Therefore classical mechanics adequately
  describes the motions of bodies much larger than
  atoms, while quantum mechanics describes the
  motions of subatomic particles and atoms as
  waves.

45
5.3 Section Quiz.

• 1. The lines in the emission spectrum for an
  element are caused by
• a) the movement of electrons from lower to higher
  energy levels.
• b) the movement of electrons from higher to lower
  energy levels.
• c) the electron configuration in the ground
  state.
• d) the electron configuration of an atom.

46
5.3 Section Quiz.

• 2. Spectral lines in a series become closer
  together as n increases because the
• a) energy levels have similar values.
• b) energy levels become farther apart.
• c) atom is approaching ground state.
• d) electrons are being emitted at a slower rate.

47
What you need to know Chapter 5

• Atomic people Democritus, Dalton, Thomson,
  Rutherford, Bohr, Schrodinger, Einstein/de
  Broglie
• Atomic orbitals (Schrodinger)energy levels,
  orbital shapes and location
• Electron configurations and orbital filling
  diagrams (Aufbau, Pauli exclusion, Hund)
• Electromagnetic radiation (visible spectrum,
  ROYGBIV, size of wavelength vs. energy, emission
  spectra
• concept of atoms absorbing a quantum of energy,
  moving to higher energy level, and emitting
  photon upon returning to lower energy level, or
  ground state see the photon in terms of color or
  light emission quanta of energy obtained by high
  temperature or high voltage
• Trial by Fire and Emission Spectra Labslight
  seen by heat and high voltage
• Vocabularyquantum mechanical model, atomic
  orbitals, electromagnetic radiation, atomic
  emission spectrum, quanta (quantum), ground
  state, photon, classical mechanics, energy
  levels, electron configuration, orbital filling,

48
Paul Dirac Carl Anderson, 1928-30
Extra stuff something to think about

• Dirac postulated that a positively charged
  electron should exist.
• Anderson discovers a particle carrying a positive
  charge of the same magnitude as an electron. He
  names it the positron. (He also wanted to
  rename the electron the negatron.

49
What is Antimatter?

• For some reason (scientists are still working on
  why), the universe is made of what we have called
  matter. The rules of matter include
• Electrons have a negative charge. Protons have a
  positive charge.
• The universe also contains what we have called
  antimatter. The rules of antimatter include
• Electrons have a positive charge. Protons have a
  negative charge.
• Antimatter is very rare
• Positrons are produced in a certain type of
  radioactive decay.
• Whenever a antimatter particle collides with its
  matter particle partner, they annihilate,
  releasing energy.
• Other anti-matter particles have been found,
  including the anti-proton and anti-neutron.

50
Murray Gell-Mann George Zweig, 1964

• Protons neutrons are made up of smaller
  particles called quarks. (Quark comes from a
  James Joyce poem.)
• There are 6 flavors of quarks now known up,
  down, top, bottom, strange, and charm.
• These were discovered using particle
  accelerators colliding particles at high speeds
  and studying the shrapnel of matter and energy.

51
History of Electric Charge

• Discovered by Ben Franklin in the 1750s in his
  famous kite experiment. Franklin coined the
  following terms to help describe his findings
• positive (to indicate an object that donates
  electric charge)
• negative (an object that accepts electric charge)
• Plus, minus, battery, conductor, and charge
• Unfortunately, Franklin decided to describe
  electricity as moving from () to (-). Later, it
  was discovered (by Thomson) that electricity is
  the motion of negative charges (electrons), which
  move from the (-) terminal on a battery to the
  () one.
• Today, physicists and physics students still use
  the convention that electricity flows from () to
  (-), even though they know that electrons flow
  the opposite way!