Electron Behavior and Electrochemistry - PowerPoint PPT Presentation

1 / 88
About This Presentation
Title:

Electron Behavior and Electrochemistry

Description:

... eight elements, and proposed an arrangement of octaves, as in music. ... Cause of Periodic Trends. Number of Protons in nucleus ... – PowerPoint PPT presentation

Number of Views:59
Avg rating:3.0/5.0
Slides: 89
Provided by: mad795
Category:

less

Transcript and Presenter's Notes

Title: 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
Write a Comment
User Comments (0)
About PowerShow.com