Title: Electron Behavior and Electrochemistry
1Electron Behavior and Electrochemistry
2Electromagnetic Radiation
- Radiant energy that exhibits wavelength-like
behavior and travels through space at the speed
of light in a vacuum. Caused by electronic
interactions.
3Electromagnetic 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.
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5Wavelength and frequency can be interconverted.
- ? c/?
- ? frequency (s?1)
- ? wavelength (m)
- c speed of light (m s?1)
6Examplep281 Zum
7Plancks 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
8Examplep283 Zum
9Energy and Mass
- Energy has mass
- E mc2
- E energy
- m mass
- c speed of light
10Energy and Mass
(Hence the dual nature of light.)
11Wavelength 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
12Examplep285 Zum
13Wavelength 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
14Atomic Spectrum of Hydrogen
- Continuous spectrum Contains all the
wavelengths of light. - Line (discrete) spectrum Contains only some of
the wavelengths of light.
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17The 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).
18The Bohr Model
- Energy Changes in the Hydrogen Atom
- ?E Efinal state ? Einitial state
19Examplep291 Zum
20Examplep293 Zum
21HOMEWORK!!
- p. 333ff 23, 24, 27, 31, 34, 37
22Polyelectronic 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
23Quantum 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
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25Heisenberg 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.
26Probability 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.
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28Quantum 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.
29Examplesp334 47,48
30p
s
f
d
31Pauli 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.
32Electron 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)
33Hunds 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.
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35Examplep314 Zum
36Valence Electrons
The electrons in the outermost principle quantum
level of an atom.
Inner electrons are called core electrons.
37Electron 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.
38Examples
39Homework
- p. 334ff 55, 60, 63, 64, 65, 66
40History 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.
41Zeff 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
42Electron 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
43Ionization Energy
- The quantity of energy required to remove an
electron from the gaseous atom or ion.
44Periodic Trends
- First ionization energy
- increases from left to right across a period
- decreases going down a group.
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46Cause 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
47Subsequent 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
48Electron 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)
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50Periodic Trends
- Atomic Radii
- decrease going from left to right across a
period - increase going down a group.
- (Lower Left Larger)
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53Cause 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
54Information 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.
55Homework!!!
- p. 335ff 67, 69, 71, 75, 82, 103, 109, 124
- Super XCR!! 128
56Electrochemistry
- 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)
57Review 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
58Half-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-
59Rules 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
60Galvanic Cell
- A device in which chemical energy is changed to
electrical energy.
61Anode and Cathode
- OXIDATION occurs at the ANODE.
- REDUCTION occurs at the CATHODE.
62Structure 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
63Cell Potential
- Cell Potential or Electromotive Force (emf) The
pull or driving force on the electrons.
64Standard 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
65Standard 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
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68Using 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
69Using 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
70Cell 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
71Homework!!
72Line 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)
73Dependence 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.
74Concentration 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
75Example 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.
76The 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
77The 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
78Example 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
79Batteries
- A battery is a galvanic cell or, more commonly, a
group of galvanic cells connected in series.
80Fuel 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-
81Corrosion
- 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.
82Rusting
- 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.
83Cathodic 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.
84Electrolysis
- 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.
85Aluminum 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.
86The 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.
87Electrolysis 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.
88Homework!!
- p. 854ff 28, 29, 47, 51, 52, 68, 71