Electron Behavior and Electrochemistry

1
Electron Behavior and Electrochemistry
2
Electromagnetic Radiation

• Radiant energy that exhibits wavelength-like
  behavior and travels through space at the speed
  of light in a vacuum. Caused by electronic
  interactions.

3
Electromagnetic Waves

• Waves have 3 primary characteristics
• 1. Wavelength distance between corresponding
  points on consecutive waves
• 2. Frequency number of waves per second that
  pass a given point in space.
• 3. Speed speed of light is 2.9979 ? 108 m/s.

4
(No Transcript)
5
Wavelength and frequency can be interconverted.

• ? c/?
• ? frequency (s?1)
• ? wavelength (m)
• c speed of light (m s?1)

6
Examplep281 Zum
7
Plancks Constant
Transfer of energy is quantized, and can only
occur in discrete units, called quanta. The
individual particles are called photons.

• ?E change in energy, in J
• h Plancks constant, 6.626 ? 10?34 J s
• ? frequency, in s?1
• ? wavelength, in m

8
Examplep283 Zum
9
Energy and Mass

• Energy has mass
• E mc2
• E energy
• m mass
• c speed of light

10
Energy and Mass
(Hence the dual nature of light.)
11
Wavelength and Mass
Louis deBroglie showed that any moving mass has
an associated wavelength de Broglies Equation

• ? wavelength, in m
• h Plancks constant, 6.626 ? 10?34 J s kg
  m2 s?1
• m mass, in kg
• v velocity, in m/s

12
Examplep285 Zum
13
Wavelength and Mass

• The concept that mass has wavelength makes the
  electron and Field Effect microscopes possible
• The electron microscope uses electrons to make a
  picture of tiny objects as they diffract or bend
  around the corners of those objects
• X-ray diffraction works in the same way to see
  atoms in crystals

14
Atomic Spectrum of Hydrogen

• Continuous spectrum Contains all the
  wavelengths of light.
• Line (discrete) spectrum Contains only some of
  the wavelengths of light.

15
(No Transcript)
16
(No Transcript)
17
The Bohr Model
The electron in a hydrogen atom moves around the
nucleus only in certain allowed circular orbits.

• E energy of the levels in the H-atom
• z nuclear charge (for H, z 1)
• n an integer
• Ground State The lowest possible energy state
  for an atom (n 1).

18
The Bohr Model

• Energy Changes in the Hydrogen Atom
• ?E Efinal state ? Einitial state

19
Examplep291 Zum
20
Examplep293 Zum
21
HOMEWORK!!

• p. 333ff 23, 24, 27, 31, 34, 37

22
Polyelectronic Atoms

• Bohrs calculations ONLY work for Hydrogen, not
  any other atom. H has only one electron, multiple
  electrons cause the difference.
• The major reason Bohrs equations dont work is
  due to the interaction of electrons with each
  otherelectron repulsion and electron shielding,
  both of which affect the pull of the nucleus on
  other electrons.
• Both of these lead to the concept of effective
  nuclear charge, which can be represented as
  Zactual core electrons
• The number of core electrons is equal to the
  atomic number of the previous Noble Gas

23
Quantum Mechanics

• First developed by Erwin Schroedinger
• Based on the wave properties of the atom
• ? wave function
• mathematical operator
• E total energy of the atom
• A specific wave function is often called an
  orbital.
• The calculations here are way beyond this course

24
(No Transcript)
25
Heisenberg Uncertainty Principle

• x position
• mv momentum
• h Plancks constant
• The more accurately we know a particles
  position, the less accurately we can know its
  momentum. In other words, we cant know where
  the electron is and how fast it is going at the
  same time.

26
Probability Distribution

• square of the wave function
• probability of finding an electron at a given
  position
• Radial probability distribution is the
  probability distribution in each spherical shell.
  This is limited to a size approximately 90 of
  the radius.

27
(No Transcript)
28
Quantum Numbers (QN)

• 1. Principal QN (n 1, 2, 3, . . .) - related
  to size and energy of the orbital.
• 2. Angular Momentum QN (l 0 to n ? 1) -
  relates to shape of the orbital.
• 3. Magnetic QN (ml l to ?l) - relates to
  orientation of the orbital in space relative to
  other orbitals.
• 4. Electron Spin QN (ms 1/2, ?1/2) - relates
  to the spin states of the electrons.

29
Examplesp334 47,48
30
p
s
f
d
31
Pauli Exclusion Principle

• In a given atom, no two electrons can have the
  same set of four quantum numbers (n, l, ml, ms).
• Therefore, an orbital can hold only two
  electrons, and they must have opposite spins. The
  reason is magnetic, since spinning electrons
  become magnets. Though they hate each other,
  they will stay in the same orbital if there is
  another attraction, magnetism, to hold them there.

32
Electron Arrangement The Aufbau Principle

• As protons are added one by one to the nucleus
  to build up the elements, electrons are similarly
  added to hydrogen-like orbitals. Electrons are
  added in order of increasing energy of those
  orbitals.
• (Electrons fill lowest energy levels first)

33
Hunds Rule

• The lowest energy configuration for an atom is
  the one having the maximum number of unpaired
  electrons allowed by the Pauli principle in a
  particular set of degenerate orbitals. This
  rule simplifies to One in each before two in
  any for orbitals of the same energy, such as p,
  d, or f orbitals.

34
(No Transcript)
35
Examplep314 Zum
36
Valence Electrons
The electrons in the outermost principle quantum
level of an atom.
Inner electrons are called core electrons.
37
Electron Arrangements

• The periodic table can be subdivided into
  subshell blocks, s, p, d, and f.
• Elements with s outer electrons are in columns 1
  and 2, the alkali and alkaline earth elements. H
  and He also belong in this block.
• Columns 3A through 8A are the p block
• The transition metals make up the d block, and
  the actinide and lanthanide rows make up the f
  block.

38
Examples
39
Homework

• p. 334ff 55, 60, 63, 64, 65, 66

40
History of the Periodic Law

• Dobereiner noticed several groups of three
  elements with similar properties, which he called
  triads.
• Newlands found that in elements he knew about,
  properties seemed to repeat every eight elements,
  and proposed an arrangement of octaves, as in
  music.
• The best effort was done by Meyer and Mendeleev
  independently, both arranged elements by weight
  and properties. Mendeleev is given ultimate
  credit due to his ability to use his result to
  predict new elements.
• Current periodic tables are based on the work of
  Moseley, who found that atomic number is more
  consistent with arrangement of properties.

41
Zeff Across the Table

• Zeff causes differences in elements across a
  period
• Outer electrons feel fewer protons than core
  electrons more protons felt greater pull on
  electrons
• Inner electrons shield outer, but not perfectly,
  due to penetration effect electrons penetrate
  to lower levels

42
Electron Interactions

• Electron subshells have increasing energies
• s lt p lt d lt f
• Electrons in a given orbital do not shield others
  in that orbital
• Electrons in lower energy orbital somewhat shield
  those in higher orbitals
• Electronic repulsion within orbitals can cause
  differences

43
Ionization Energy

• The quantity of energy required to remove an
  electron from the gaseous atom or ion.

44
Periodic Trends

• First ionization energy
• increases from left to right across a period
• decreases going down a group.

45
(No Transcript)
46
Cause of Periodic Failures

• Borons Lesser Ionization
• Caused by partial shielding by the 2s electrons
• Not very large amount, but explains the gap
• Oxygens Lesser Ionization
• Caused by electron repulsion in the 2p orbital
• More energy to the 2nd electron, easier to remove

47
Subsequent Ionization Energies

• Once an element has lost one electron, taking
  away another is harder, due to increased positive
  charge of the atom
• This is especially noticeable in trying to take
  core electrons away, which are held much tighter
  than valence electrons

48
Electron Affinity

• The energy change associated with the addition
  of an electron to a gaseous atom. Trend similar
  to ionization energy.
• X(g) e? ? X?(g)

49
(No Transcript)
50
Periodic Trends

• Atomic Radii
• decrease going from left to right across a
  period
• increase going down a group.
• (Lower Left Larger)

51
(No Transcript)
52
(No Transcript)
53
Cause of Periodic Trends

• Number of Protons in nucleus
• More protons, more pull on the outer electrons
• Shielding effect of Core electrons
• Offsets the pull of the protons
• Equal to Atomic Number minus inner shell electrons

54
Information Contained in the Periodic Table

• 1. Each group member has the same valence
  electron configuration (these electrons primarily
  determine an atoms chemistry).
• 2. The electron configuration of any
  representative element.
• 3. Certain groups have special names (alkali
  metals, halogens, etc).
• 4. Metals and nonmetals are characterized by
  their chemical and physical properties.

55
Homework!!!

• p. 335ff 67, 69, 71, 75, 82, 103, 109, 124
• Super XCR!! 128

56
Electrochemistry

• The study of the interchange of chemical and
  electrical energy by the transferring of
  electrons between substances
• Electrochemistry involves the principles of
  oxidation and reduction in both galvanic action
  (producing electricity from chemical changes) and
  electrolytic action (producing chemical changes
  with electric current)

57
Review of Terms

• oxidation-reduction (redox) reaction involves a
  transfer of electrons from the reducing agent to
  the oxidizing agent.
• oxidation loss of electrons
• reduction gain of electrons

58
Half-Reactions

• The overall reaction is split into two
  half-reactions, one involving oxidation and one
  reduction.
• 8H MnO4- 5Fe2 Mn2 5Fe3 4H2O
• Reduction 8H MnO4- 5e- Mn2 4H2O
• Oxidation 5Fe2 5Fe3 5e-

59
Rules for Assigning Oxidation States

• 1. Oxidation state of an atom in an element 0
• 2. Oxidation state of monatomic ion charge
• 3. Oxygen ?2 in covalent compounds (except in
  peroxides where it ?1)
• 4. H 1 in covalent compounds pg159
• 5. Fluorine ?1 in compounds
• Sum of oxidation states 0 in compounds
• Sum of oxidation states charge of a polyatomic
  ion

60
Galvanic Cell

• A device in which chemical energy is changed to
  electrical energy.

61
Anode and Cathode

• OXIDATION occurs at the ANODE.
• REDUCTION occurs at the CATHODE.

62
Structure of Galvanic Cells

• The oxidation and reduction can only continue if
  electrons can flow between cells.
• This requires a salt bridge or porous disk to
  allow for balance of charge between the two
  half-reactions
• When connected by wires across a device, current
  flows.

Porous disk
63
Cell Potential

• Cell Potential or Electromotive Force (emf) The
  pull or driving force on the electrons.

64
Standard Reduction Potentials

• The E values corresponding to reduction
  half-reactions with all solutes at 1M and all
  gases at 1 atm.
• Cu2 2e- Cu E 0.34 V vs. SHE
• SO42- 4H 2e- H2SO3 H2O
• E 0.20 V vs. SHE

65
Standard Reduction Potentials

• Half cells provide different abilities to push or
  pull electrons. Since reduction has to do with
  gain of electrons, it is considered as pulling
  them, and given the positive sign.
• Cells are measured in regard to how strongly they
  pull electrons compared to hydrogen.
• Some cells have a negative pull, and thus a
  negative reduction potential

66
(No Transcript)
67
(No Transcript)
68
Using Reduction Potentials

• Sample 17-1A Consider a galvanic cell based on
  the reaction
• Al3(aq) Mg(s)? Al(s) Mg2(aq)
• Use the appropriate half-reactions and calculate
  E o for the cell

69
Using Reduction Potentials

• 17-1B Consider a galvanic cell based on the
  reaction
• MnO4-(aq) H(aq) ClO3-(aq)? ClO4-(aq)
  Mn2(aq)H2O(l)
• Use the appropriate half-reactions and calculate
  E o for the cell

70
Cell Voltages and Spontaneity

• A reaction which proceeds without any outside
  help (heat, electricity, etc) is called
  spontaneous. Spontaneous reactions are not
  necessarily fast, but they do proceed.
• A spontaneous redox reaction will have a positive
  E , a non-spontaneous one will have a negative
  E

71
Homework!!

• p. 853 17, 18, 19, 20

72
Line Notation

• Half cells are often referred to by special
  notation
• From sample 17-1A, the reaction there would be
  shown as
• Mg(s)Mg2(aq)Al3(aq)Al(s)
• Phase differences Salt bridge

Example 17-1Bs components are all ions, so no
phase difference is present, and the electrodes
must be some inert metal, usually Pt. This cell
would be designated Pt(s)ClO3-(aq),ClO4-(aq)M
nO4-(aq), Mn2(aq)Pt(s)
73
Dependence on Concentration

• Cells have variation of their potential depending
  on concentration of reactants and products
• If reactants are increased, the cell will make
  slightly higher potential, if products are
  increased, the potential will be decreased.
• It follows then that the cell will lose potential
  as reactants turn into products, thus why
  batteries run down.

74
Concentration Cell

• . . . a cell in which both compartments have the
  same components but at different concentrations.
• Current flows due to the imbalance of the
  concentrations

Ag off
Ag on
75
Example 17.6
Determine the direction of electron flow and
designate anode and cathode for the cell
represented here
0.01 M Fe3
0.1 M Fe3
The 0.01 M solution is too weak compared to the
other, so Fe will lose electrons and become Fe3
in the left cell. The electrons will flow to the
other cell where they will combine with Fe3 ions
to form Fe, thus reducing the concentration in
that cell. The left side has oxidation taking
place, so it is the anode, the right side has
reduction, so it is the cathode.
76
The Nernst Equation

• We can calculate the potential of a cell in which
  some or all of the components are not in their
  standard states.
• Where Q the ratio of aqueous product
  concentration to aqueous reactant concentration

E E o
77
The Nernst Equation

• Since the ratio of reactants to products changes
  with time, the value of Q increases, thus
  increasing the negative factor and decreasing the
  E for the cell

E E o
78
Example 17.5

• For the cell reaction
• 2Al(s) 3 Mn2(aq)? 2 Al3(aq) 3 Mn(s) E
  o0.48 V
• Use the Nernst equation to predict if Ecell is
  larger or smaller than E o
• (a) Al3 2.0 M , Mn21.0 M (b)Al3
  1.0 M , Mn23.0 M

79
Batteries

• A battery is a galvanic cell or, more commonly, a
  group of galvanic cells connected in series.

80
Fuel Cells

• . . . galvanic cells for which the reactants are
  continuously supplied.
• 2H2(g) O2(g) 2H2O(l)
• anode 2H2 4OH- 4H2O 4e-
• cathode 4e- O2 2H2O 4OH-

81
Corrosion

• The concept of corrosion, such as rusting of
  iron, is considered an oxidation/reduction
  process. Many metals create a thin layer of
  protective oxide on their surfaces, such as zinc
  and aluminum. Iron also forms such a coating,
  but it does not adhere to the metal, and flakes
  off, exposing more iron to the air.
• Some metals, such as copper, gold, silver and
  platinum, are relatively difficult to oxidize.
  They do not form oxides, and hardly react at all.
  These are often called noble metals.

82
Rusting

• The corrosion if iron is an important reaction
  which we want to inhibit. First, water and
  oxygen must be kept away, then all rust particles
  must be removed, since they aid in the
  destruction of more iron, as shown in the
  picture.
• The most common ways of avoiding rusting are
  plating with a less reactive metal, coating with
  a sacrificial metal, and alloying.

83
Cathodic Protection

• Steel can be protected by coating with zinc, a
  process known as galvanizing. The zinc takes
  over as the anode and becomes oxidized to ZnO,
  which protects the sheet metal.
• Buried pipes and water heaters are protected by
  attaching a more active metal, such as aluminum
  or magnesium to the steel. The active metal
  corrodes instead of the steel, because it becomes
  the anode, whereas the steel is then the cathode.

84
Electrolysis

• If an electrochemical cell ends up with a
  negative E , it may be forced to work by adding
  electric current, forcing current through a cell
  to produce that chemical change.
• Electrolysis can be used to break up water, cause
  elements to be plated to other metals, and to
  purify metals from impure materials.

85
Aluminum is prepared by the Hall-Heroult process,
by which aluminum oxide (bauxite) has an electric
current passed through it.
This process has resulted in the price of
aluminum dropping from as much as 10,000 per
pound to as little as 0.30 per pound since the
development of the process.
86
The process of electroplating involves sending
electric current into an object to be plated
which is immersed in a solution of the metal
which is to be plated onto it. A bar of the pure
metal is the anode, which sends its ions into
solution as the current flows. The ions then
plate onto the item as long as current is flowing
in the cell.
87
Electrolysis of Various Compounds
If molten sodium chloride is electrolyzed, both
sodium and chlorine can be obtained. This is the
process by which sodium is made for use in
chemical labs. If aqueous NaCl is used instead,
sodium is not obtained, since it is less likely
to be reduced than hydrogen from water.
The aqueous process makes chlorine gas at the
anode, hydrogen gas and hydroxide at the cathode.
This process is the primary process for making
sodium hydroxide, which is used in industry.
88
Homework!!

• p. 854ff 28, 29, 47, 51, 52, 68, 71