Chapter 4:Arrangement of Electrons in Atoms

1
Chapter 4Arrangement of Electrons in Atoms
Chapter 4Arrangement of Electrons in
Atoms Date Started 11-10-2008 Date
Completed____________ Period ____
Your Name
2
Chapter 4

• Properties of Light
• Electromagnetic radiation is a form of energy
  that exhibits wavelike behavior as it travels
  through space.
• Visible light and all other forms of
  electromagnetic radiation make up the
  electromagnetic spectrum.
• All forms of electromagnetic radiation travel (in
  a vacuum) at a fixed or constant speed of 300,000
  km /second or 186,000 miles/second. This is
  called the speed of light.
• The wave motion is periodic or has a repetitive
  nature and is described by wavelength and
  frequency.
• Wavelength is the distance between corresponding
  points on adjacent waves and is measured in
  meters.
• Frequency is the number of waves that pass a
  given point in a specific time, usually in one
  second. It is measured in hertz (Hz) and is
  (waves) cycles per second.
• The period of a wave is the time it takes a wave
  to make one complete cycle.

3
Wavelength and Frequency
4
The Photoelectric effect

• 1.  When certain metals are exposed to a
  specific wavelength of light, these metals give
  off electrons. This is known as the
  photoelectric effect.
• 2.     Max Planck proposed that objects emit
  energy in small, specific amounts called quanta.
  A quantum is t the minimum amount or quantity of
  energy that can be lost or gained by an atom.
• 3.     Because of the photoelectric effect, the
  wave theory of light could not explain this and
  led Albert Einstein to propose that light had
  the property of particles.

5
Photos of Photoelectric effect
6
The Photoelectric effect

• 4.  Einstein proposed the concept of the photon,
  which is a particle of light, carries a quantum
  of energy, which has no rest mass.
• 5.  Einstein proposed in his theory that when
  matter absorbs photons, the photon strikes an
  electron, the electron is knocked loose from its
  orbit.
• 6.  The ground state of an atom is the lowest
  energy state where electrons normally are found.
• 7.  The excited state of an atom is where the
  atom has absorbed energy (photons) and has a
  higher potential energy compared to the ground
  state.

7
Speed of Light Formula

• c l n
• Where
• c speed of light (300 000 km/sec)
• l wavelength (in kilometers)
• n frequency (in Hertz)

8
   ENERGY AND LIGHT

• 1.  All light has energy associated with it to
  perform work.
•  2. The amount of energy is based on the
  frequency or wavelength of light.
• 3. The amount of energy is directly
  proportional to the frequency times Planck's
  constant.
• E h v
• where E energy in joules of a photon
  of radiation
• h 6.6 x 10-34 joules per hertz
• v (f ) frequency in hertz or cycles per
  second.

9
Continuous Spectrum IR to UV light

•  
• R O Y G
  B I V
• E r e
  r 1 n i
• D a l
  e u d o
• n l
  e e i l
• g o n
  g e
• e w o
  t

Low Energy High Energy Low Frequency High
Frequency
10
Three Types of Light Producing Interactions

• 1. Atom to atom interaction through
  chemical reactions.
• 2.    Electron movement from lower energy
  levels to higher energy levels and back to their
  original ground state.
•    3. Nuclear particle interactions through
  radioactive decay.

11
The Continuous Spectrum

• Spectroscopy the analysis of light being
  emitted or received from an object using a
  spectroscope, in order to determine it's atomic
  composition.
• Light is produced due to energy being given
  by an atom with three types of interactions,
  which can produce light.
• The visible spectrum is a very small and
  narrow range of frequencies.    
•  

12
(No Transcript)
13
TYPES of Light Spectra
1.   Continuous Spectra all forms of visible and
invisible electromagnetic radiation. 2.  
Absorption Spectra most of light in this pattern
is shows absorption of almost all light by an
atom. 3.   Emission Spectra only light of
certain wavelengths are given off by atoms.
14
How are spectra produced?

• An Absorption Spectrum is produced when electrons
  move to higher energy levels.
• The greater the number of electrons moving up to
  higher energy levels, (excited state) the
  greater the absorption bands that are
  present.
• An Emission Spectrum is produced when electrons
  that have moved to higher energy levels,
  return to their original energy level.
• The color of the light emitted is based on how
  much energy the electrons must release, in
  order to return to it's ground state.
• Red light has a lower frequency than blue light,
  and therefore, has less energy than blue
  light. (Planck's formula for finding
  energy).
•  

15
Red Light
16
Blue Light
17
Green Light
18
White Light
19
Black Light
20
Emission Spectra
21
Quick Check

• What are two components of light?
• What do you get when these two components are
  multiplied together?
• Using the color spectrum of the rainbow, what
  color of light has the lowest energy?
• Using the color spectrum of the rainbow, what
  color of light has the highest energy?
• What is the name of the concept where when a
  specific color of light hits a metal, and it
  causes electrons to be released from their
  orbits?

22
Answers to Quick Check

• Light has wavelength and frequency. Wavelength
  is the length of the wave in meters, while
  frequency is the number of waves per second.
• The product of wavelength and frequency is the
  speed of a light. (c)
• Red light has the lowest energy because it has
  the lowest frequency of visible light.
• Violet light has the highest energy because it
  has the highest frequency of visible light.
• The photoelectric effect, proposed by Albert
  Einstein

23
Hydrogen-Atom Line Emission

• When hydrogen atoms are excited with energy,
  spectral lines of light are produced.
• These spectral lines represent the energy lost by
  atoms that were in the excited state.
• There are three series of spectral lines that
  hydrogen atoms can produce when excited.
• They are
• Balmer series (5 lines) in the visible spectrum
• Lyman Series (6 lines) in the UV spectrum
• Paschen Series (4 lines) in the IR spectrum

24
Hydrogen Line Emission
Balmer Series for Hydrogen Gas
25
Hydrogen Line Emission Chart
High Energy Low Energy
26
Modern Atomic TheoryBohr Model (Hydrogen)

• Electrons orbit the nucleus at different
  distances.
• The distance and the type of orbit that an
  electron has is determined by the Kinetic energy
  the electron possesses.
• Electrons can move to other energy levels only if
  they gain or lose energy.
• The greater the distance an electron is from the
  nucleus, the higher the amount of kinetic energy
  that is present in the electron.
• Electrons fill orbits or energy levels based on
  their kinetic energy and do so in an orderly
  manner (AUFBAU)

27
Modern Atomic TheoryBohr Model page 2

• There are seven MAJOR energy levels that
  electrons can occupy or move to and from.
• The energy levels are like ladders, the can only
  hold a certain number of electrons (Area of a
  circle).
• When an electron falls from a higher energy level
  to a lower energy level, the electron gives off
  this energy in the form of light (Emission
  spectrum)
• When an electron jump from a lower energy level
  to a higher energy level, the electron must
  absorb this energy in the form of light
  (Absorption spectrum)

28
Ground and Excited States

• Ground State electron
• Excited State electron

29
Quantum Model of an Atom

• Electrons show a dual wave-particle behavior just
  as light energy.
• In 1924, Louis DeBroglie used diffraction and
  interference patterns of light and compared these
  to the electron patterns in excited crystals.
• In 1927, Werner Heisenberg proposed his
  uncertainty principle to explain the location of
  electrons and their speed.
• Heisenberg Uncertainty Principle states that it
  is impossible to determine simultaneously both
  the position and velocity of an electron or any
  particle.
• Quantum theory describes mathematically the wave
  properties of electrons, and the type of orbit
  that electrons demonstrate.
• Quantum numbers specify the properties of
  electron orbits in terms of their orbital
  location, shape and capacity.

30
Diffraction and Interference
31
Schrodingers Quantum Theory

• 1.  Electrons are found above the nucleus in 3D
  regions of space called orbitals.
• 2.  These orbital bands are highly probable
  regions where the electrons can be found. Only
  two electrons can occupy one orbital.
• 3.   Electrons can shift to higher energy levels
  when they have gained kinetic energy.
  Eventually these electrons lose this energy and
  return to their original orbital.
• 4.   Line spectra are produced when electrons
  drop from higher energy levels (excited state)
  to lower energy levels.
• 5.    Quantum numbers specify or note the
  properties of atomic orbitals and their
  electronsgt There are four quantum numbers.

32
Four Quantum Numbers

• Principle Quantum Number The symbol (n)
  represents the number of energy levels for
  electrons in an atom. These are whole numbers
  from 1 to 7.
• Angular (orbital/momentum) Quantum Number The
  symbol is (l) and represents the shape of the
  electron orbitals and are called sublevels. For
  this quantum number l is n-1 and is noted by the
  letterss, p, d, and f.
• Magnetic Quantum Number The symbol is m and
  represents the orientation of the orbital around
  the nucleus in 3D space.
• Spin Quantum Number The symbol is (½ ) or (-½
  ) and represents the direction of spin of an
  electron in oreintation in the orbital around the
  nucleus. (cw-clockwise or ccw-counter-clockwise)

33
Additional Quantum Principles

• Heisenberg Uncertainty Principle states that it
  is impossible to determine simultaneously both
  the position and velocity of an electron or any
  particle.
• Pauli exclusion principle No two electrons in
  the same atom can have the same set of quantum
  numbers.
• Hunds Rule Orbitals of equal energy are each
  occupied by one electron before any orbital is
  occupied by a second electron, and all electrons
  in singly occupied orbitals have the same spin.

34
Quantum Number Chart
See page 5 in note packet
35
Electron Orbital Shapes

• s orbital is spherical in shape
• p orbital is peanut or lobed in shape
• d orbital is dumbbell in shape.

36
p - Orbital Orientations
37
d - Orbital Orientations
38
f- orbital Orientations
39
Quick Check 2

• What scientist developed the uncertainty
  principle?
• What is called the highly probable region that
  electrons can be found in the atom?
• How many different types of sublevels are there
  in an atom?
• What are the four letters used to show the
  different types of sublevels?
• What does the equation 2n2 ?

40
Quick Check 2

• Heisenberg
• An orbital
• Four different types of sublevels
• s spherical shape
• p peanut shaped
• d dumbbell shaped
• f fractionated shapes
• 2n2 the number of electrons in an
  energy level?

41
Energy overlaps of Orbitals
42
Aufbau Handout
43
Writing Electron Configurations

• 1. Determine the total number of
  electrons from the atomic number of the element.
• 2. Use Aufbau sheet for the order of
  electron fill into their orbitals.
• 3. Keep track of number of electrons
  used.
• 4. The number in front of the orbital
  letter is the energy level the orbital.
• 5. The s and p orbital electrons in the
  outer most energy level from the nucleus is
  called the valence shell.

44
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
45
Electron dot or Lewis Dot

• Chemical symbol is written for the element.
• Determine the number of valence electrons.
  (outermost s and p sublevel)
• Place dots diagonally around symbol using an
  imaginary box. (maximum of2 dots per side)

46
Orbital Notation

• Do the electron configuration of the element
• Place one line under s sublevels.
• Place three lines under p sublevels.
• Place five lines under d sublevels.
• Place seven lines under f sublevels.
• Use Hunds Rule to fill orbitals on each sublevel.

47
3 Expressions of Electron Shell Filling of Ca

• Electron configuration 1s22s22p63s23p64s2
• Lewis or Electron Dot Ca
• Orbital Notation
• __ __ __ __ __ __ __ __ __
  __
• 1s2 2s2 2p6
  3s2 3p6 4s2

48
Electron shell filling-Orbital Notation-
Electron Dot

• S (16)
• 1s2 2s2 2p6 3s2 3p4
• Sn (50)1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5s2 4d10 5p2

49
Electron shell filling-Orbital Notation-
Electron Dot

• Ni (28)
• 1s2 2s2 2p6 3s2 3p6 4s2 3d8
• Br (35)
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

50
Electron Configuration

• 1 e- - H - 1s1
• 2 e- - He - 1s2 (filled)
• 3 e- - Li - 1s2 2s1
• 4 e- - Be - 1s2 2s2
• 5 e- - B - 1s2 2s2 2p1
• 6 e- - C - 1s2 2s2 2p2
• 7 e- - N - 1s2 2s2 2p3
• 8 e- - O - 1s2 2s2 2p4
• 9 e- - F - 1s2 2s2 2p5
• 10 e- - Ne - 1s2 2s2 2p6 (filled)

51
Electron Configuration

• 11 e- - Na - 1s2 2s2 2p6 3s1
• 12 e- - Mg - 1s2 2s2 2p6 3s2
• 13 e- - Al - 1s2 2s2 2p6 3s2 3p1
• 14 e- - Si - 1s2 2s2 2p6 3s2 3p2
• 15 e- - P - 1s2 2s2 2p6 3s2 3p3
• 16 e- - S - 1s2 2s2 2p6 3s2 3p4
• 17 e- - Cl- 1s2 2s2 2p6 3s2 3p5
• 18 e- - Ar- 1s2 2s2 2p6 3s2 3p6 (filled)
• 19 e- - K- 1s2 2s2 2p6 3s2 3p6 4s1
• 20 e- - Ca - 1s2 2s2 2p6 3s2 3p6 4s2

52
Electron Configuration

• 21 e- - Sc - 1s2 2s2 2p6 3s2 3p6 4s23d1
• 22 e- - Ti - 1s2 2s2 2p6 3s2 3p6 4s23d2
• 23 e- - V - 1s2 2s2 2p6 3s2 3p6 4s23d3
• 24 e- - Cr - 1s2 2s2 2p6 3s2 3p6 4s13d5
• 25 e- - Mn- 1s2 2s2 2p6 3s2 3p6 4s23d5
• 26 e- - Fe - 1s2 2s2 2p6 3s2 3p6 4s23d6
• 27 e- - Co - 1s2 2s2 2p6 3s2 3p6 4s23d7
• 28 e- - Ni - 1s2 2s2 2p6 3s2 3p6 4s23d8
• 29 e- - Cu- 1s2 2s2 2p6 3s2 3p6 4s13d10
• 30 e- - Zn - 1s2 2s2 2p6 3s2 3p6 4s23d10

53
Electron Configuration

• 31 e- - Ga - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
• 32 e- - Ge - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
• 33 e- - As - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
• 34 e- - Se - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
• 35 e- - Br - 1s2 2s2 2p6 3s2 3p6 4s23d10 4p5
• 36 e- - Kr - 1s2 2s2 2p6 3s2 3p6 4s23d10 4p6
  (filled)
• 37 e- - Rb - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5s1
• 38 e- - Sr - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5s2
• 39 e- - Y - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5s2 4d1
• 40 e- - Zr - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5s2 4d2

54
Period 2Electron Configurations
55
Period 3Electron Configurations
56
Period 4Electron Configurations
57
Period 5Electron Configurations
58
New Sublevels in AufBau
59
3 Classes of Magnetic Materials

• Ferromagnetic materials highly magnetic, and
  retain magnetic fields these materials have a
  large number of unpaired d sublevel electrons.
  (iron, cobalt, nickel)
• Paramagnetic materials very weakly magnetic
  materialsd sublevel electrons are completely
  paired. (Aluminum)
• Diamagnetic materials atoms which have no net
  magnetic moments all the d orbital shells are
  filled and/or there are no unpaired d orbital
  electrons. When exposed to a magnetic field, a
  negative magnetization is produced and thus the
  susceptibility is negative (carbon, bismuth,
  water)

60
Magnetism and orbital pairs
61
d sublevel electrons and magnetic domains
62
Diamagnetism and LevitationVideos

• Flying frog
• Floating Grasshopper
• Floating Strawberry
• Water drop
• Tomato

63
(No Transcript)