Download Chemistry 30 unit two

Chemistry 30 unit two
http://www.youtube.com/watch?v=JJudWHG-gKM
 You will explain the nature of oxidation-reduction reactions
 You will apply the principles of oxidation-reduction to electrochemical cells.
Electrochemistry:
 You will define oxidation and reduction operationally and theoretically
Definitions:
Oxidation-Reduction Reactions: The chemical changes that occur when
electrons are transferred between reactants. Oxidation reactions are the
principal source of energy on Earth.
Eg.
First Definitions:
Oxidation: Elements and compounds combine with oxygen.
Reduction: Loss of oxygen from a compound.
-A metal oxide is reduced to the metal (smaller volume-reduced)
*oxidation and reduction occur simultaneously.
Modified Definitions: (Electron transfer in redox reactions)
Oxidation
-loss of electrons
or gain of oxygen
Reduction
-gain of electrons or
loss of oxygen
LEO the lion says GER
LEO =loss of electrons in oxidation
GER =gain of electrons in reduction
Eg.
Reducing Agent: The reactant in a Redox reaction that donates electrons. It
promotes reduction and is itself oxidized. Eg.
Oxidizing Agent: The reactant in a Redox reaction that accepts electrons.
An oxidizing agent causes something else to be oxidized and is itself
reduced. Eg.
Eg. 1 Mg + S  MgS
Eg. 2 AgNO3 + Cu
Eg. 3
2H2 +O2 2H2O
With carbon(covalent) compounds – other definitions also apply:
Oxidation: addition of oxygen or removal of hydrogen
Reduction; removal of oxygen or addition of hydrogen
SUMMARY:
Oxidation
-complete loss of e- (ionic reaction)
-shift of e-away from an atom in a covalent
bond
-gain of oxygen
-loss of hydrogen by a covalent cmpd
-an increase in oxidation number
Reduction
-complete gain of e- (ionic rxn)
-shift of e- toward an atom in a
covalent bond
-loss of oxygen
-gain of hydrogen by a covalent
compound
-a decrease in oxidation number
Assigning oxidation numbers:
Oxidation number changes:
An increase in the oxidation number of an atom = oxidation
A decrease in the oxidation number of an atom = reduction
*oxidation numbers allow chemists to keep track of e- during the chemical
rxns. However they may not represent the actual charges on atoms. (only
on simple ions in ionic substances do they represent charge!)
IDENTIFYING REDOX REACTIONS –

You will differentiate between redox reactions and other reactions, using halfreactions and/or oxidation numbers
An equation in which the oxidation number of at least two elements change
is a redox reaction.
TYPES OF REACTIONS
TYPICALLY REDOX
TYPICALLY NON-REDOX
SINGLE REPLACEMENT
COMBUSTION
SIMPLE COMPOSITION
SIMPLE DECOMPOSITION
DOUBLE REPLACEMENT
ACID-BASE
http://www.youtube.com/watch?v=UTfMrx7275w
BALANCING REDOX EQUATIONS
 You will define oxidizing agent, reducing agent, oxidation number, half-reaction,
disproportionation
 You will write and balance equations for redox reactions in acidic and neutral
solutions by
o using half-reaction equations obtained from a standard reduction potential
table
o developing simple half-reaction equations from information provided
about redox changes
o assigning oxidation numbers, where appropriate, to the species undergoing
chemical change
Methods: 1. Balance like any other reaction only making sure charge is
also conserved.
2. Changing oxidation number method
3. Half reaction method
*total number of electrons gained in reduction must equal the total number
of electrons lost in oxidation.
2. CHANGING OXIDATION NUMBER METHOD
A redox equation is balanced by comparing the increases and decreases in
oxidation numbers.
Steps: 1. Assign oxidation numbers to all of the atoms in the equation.
2. Identify which atoms are oxidized and which are reduced.
3. Use a line to connect the atoms that undergo oxidation and those
that undergo reduction. Write the oxidation number change at the
midpoint of each line (pay attention to subscripts as you do this).
4. Make the total increase in oxidation number equal to the total
decrease in oxidation number by using appropriate coefficients.
5. For acidic conditions; you can add H2O to balance oxygen and
H+ to balance hydrogen atoms and charge.
6. Balance the remainder of the equation for any other atoms and
check to see that total charges balance.
Disproportionation:
When one species is both oxidized and reduced. Occurs when a substance
can act as both an oxidizing agent AND a reducing agent (gain or lose
electrons)
Eg. 2H2O2 2H2O + O2 (oxygen is both oxidized and reduced)
3. USING HALF REACTIONS TO BALANCE/Predicting rxns
HALF-REACTION: An equation showing either the reduction or the
oxidation of a species in a redox reaction.
Steps: 1. Write the ‘skeleton’ (unbalanced) equation in total ionic form separate ionic aqueous substances and strong acids into their ions
(as they exist in solution)
2. Using oxidation numbers pick out the 2 substances that are
changing oxidation # (undergoing oxidation/reduction).
3. Write separate half-reactions for the oxidation and reduction
processes
4. Balance the atoms in each half reaction that are changing oxidation
#s.
5. Add sufficient electrons to one side of each half-rxn to balance the
oxidation #s.
6. If it’s in acidic conditions, you can add H2O to balance oxygen
and H+ to balance hydrogen. *You now have balanced halfreactions.
7. Multiply each half-rxn by an appropriate number to make the
electron changes equal in both half-rxns.
8. Cross out substances that appear on both sides of the equation and
then add the half-reactions.
9. The final equation should be checked to be sure that atoms are
conserved, charge is conserved, and all electrons have cancelled.
(You end up with a net ionic equation)
Spontaneity Rule
*** Spontaneous Reaction: Oxidizing agent is above the Reducing Agent in
a table of reduction half-reactions. (OA above RA) Occurs naturally without
added energy.
***Non-spontaneous Reaction: Oxidizing agent is below the Reducing
Agent in a table of reduction half-reactions. (RA above OA) Does not occur
without adding energy.
TABLE BUILDING
 You will compare the relative strengths of oxidizing and reducing agents, using
empirical data
 You will design an experiment to determine the reactivity of various metals
 You will use a standard reduction potential table as a tool when considering the
spontaneity of redox reactions and their products
**Work with one equation at a time.
1. Identify the OA and RA for the first equation or set of reactants.
2. Arrange them in 2 columns( OA on the left, RA on the right) using the
spontaneity rule.
3. Complete the half reaction for each substance in the new table.
4. Repeat for the next equation until the table is complete.
PREDICTING REDOX REACTIONS
 You will identify electron transfer, oxidizing agents and reducing agents in redox
reactions that occur in everyday life, in both living systems (e.g., cellular
respiration, photosynthesis) and nonliving systems; i.e., corrosion
1. List all species present, as they exist (ionic aqueous substances as
ions, STRONG acids as ions), and identify each as a possible
oxidizing agent, reducing agent, or both.
Key words in a question:
‘Acidic’ – include H+ as a reactant
‘basic’ – include OH- as a reactant
aqueous or solutions – include H2O as a reactant
2. Choose the strongest oxidizing agent (highest on the left) as indicated
in the table of redox half-reactions, and write the reduction halfreaction equation. Include the E value if asked.
3. Choose the strongest reducing agent (lowest on the right) as indicated
in the table of redox half reactions, and write the oxidation halfreaction equation. (Equation is written in reverse to the one in the
table) Include the E if asked (reverse the sign since the equation is
reversed).
4. Balance the number of electrons lost and gained in the half reaction
equations by multiplying one or both equations by a coefficient. Add
the two half reactions equations to obtain the net ionic equation. Add
E values if required.
5. predict whether the net ionic equation represents a spontaneous or
non-spontaneous redox reaction using the spontaneity rule.
REDOX STOICHIOMETRY/TITRATIONS
 You will perform calculations to determine quantities of substances involved in
redox titrations.
 You will select and correctly use the appropriate equipment to perform a redox
titration experiment
In a titration, one reagent (titrant) is slowly added to another (sample) until
an abrupt change in a solution property (endpoint) occurs. In a redox, the
endpoint is usually a color change. Permanganate and dichromate ions in
acidic solution are commonly used because they are strong oxidizing agents
and undergo a color change. The color change of the sample is due to
excess (unreacted) ions of the titrant and indicates that the reactant in the
sample is completely used up.
The volume of the titrant used when the endpoint is reached is called the
equivalence point volume. Stoichiometric quantities have reacted.
A primary standard is the chemical that can be used directly to prepare a
standard solution (solution of known concentration).
CELLS AND BATTERIES
 You will define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte,
external circuit, power supply, voltaic cell and electrolytic cell
Volta’s Battery
A version of Volta’s first
battery. Each beaker
contains two different metals,
copper and zinc, in an
electrolyte, salt water. A
series of beakers forms a
series of cells (a battery)
whose total voltage is the
sum of the individual voltages
of all cells.
Volta’s revised cell design, simpler than the
first, consisted of a sandwich of two metals
separated by paper soaked in salt water (the
electrolyte). A cell consisted of a layer of zinc
metal separated from a layer of copper metal
by the brine-soaked paper. A large pile of cells
could be constructed to give more electrical
energy per unit charge.
-A ‘cell’ consists of an electrolyte – aqueous conductor, and two
electrodes – solid conductors. One of the electrodes provides the surface for
reduction and is called the CATHODE (CPR). The other electrode provides
the surface for oxidation and is called the ANODE (ANO). Electrons flow
through the external circuit from anode to cathode.
-more than one cell, connected together from cathode to anode (or
vice versa), is called a BATTERY.
-VOLTMETER – measures Electric Potential Difference (voltage V) or the
change in energy per charge as it moves from one position to another. The
voltage of a cell is a measure of the change in energy per charge from one
electrode to the other. Since voltage is a ratio of energy to charge, it is not
dependent on the size of the cell. (AA,B,C, D cells are all 1.5V). The
larger the cell, the more energy it can store and the more it can deliver at one
time. Therefore, the voltage of a cell depends on the chemical composition
of the reactants in the cell.
-AMMETER – measures the Electric Current (amperes A) which is
the rate of flow of charge past a point in the circuit. The larger the cell (eg.
D), the greater the current that can be produced by the cell.
-The Power (Watts – W) of a cell is determined by the combination of
both current and voltage (P=IV)
Electrochemical Cells – convert electricity to chemical
energy(electrolytic) or chemical energy to electricity (voltaic)
 You will identify the similarities and differences between the operation of a
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voltaic cell and that of an electrolytic cell
You will predict and write the half-reaction equation that occurs at each electrode
in an electrochemical cell
construct and observe electrochemical cells
You will design an experiment, including a labelled diagram, to test predictions
regarding spontaneity, products and the standard cell potential for reactions
occurring in electrochemical cells
You will identify the products of electrochemical cells
You will compare predictions with observations of electrochemical cells
You will identify the limitations of data collected on an electrochemical cell
You will explain the discrepancies between the theoretical and actual cell
potential
VOLTAIC CELLS
To describe what happens in each part of the cell, chemists separate the half
reactions by separating the two sides of a cell. A half-cell consists of one
electrode and one electrolyte. Two half-cells are required and electrolytes
are connected by a porous boundary. The boundary permits ions to move
into and out of it through tiny openings in the cotton plugs of a salt bridge or
in the walls of a porcelain cup.
Shorthand or ‘cell notation’ is:
A single line indicates a phase boundary (solid from liquid), a double line
represents a physical boundary (salt bridge or porous cup). It is common
practice to write the anode on the left, cathode on the right.
Eg. Silver /copper voltaic cell
Theory: Salt bridge balances charges to allow for continous flow: otherwise
a build up of charge would occur on each side and prohibit further reactions.
Cations to cathode/Anions to the anode!!!
Inert electrodes:
An inert electrode is one that does not react but provides a surface for
reaction. It must be a conductor but it also must not dissolve or be very
reactive as you are generally trying to cause something in the solution to
react instead. Common inert electrodes are carbon rod and platinum foil.
Carbon is cheaper!
Eg. Copper half cell with carbon and acidified dichromate
ELECTRICAL POTENTIAL EO
 You will predict the spontaneity or nonspontaneity of a redox reaction, based on
standard reduction/cell potentials, and the relative positions of the half-reaction
equations on a standard reduction potential table. You will compare their
predictions to experimental results.
 You will explain that the values of standard reduction potential are all relative to 0
volts, as set for the hydrogen electrode at standard conditions
 You will calculate the standard cell potential for electrochemical cells
The standard reduction potential (from your data booklet) of a half-reaction
is a comparative measure showing the spontaneity of rxns with respect to
hydrogen. Hydrogen was used to compare the other half-reactions to in
order to obtain the values since both an OA and RA are required for a
reaction. (see pg 628 Hydrogen half-cell set-up)
Changing the substance of comparison (hydrogen) to another substance will
change the value for the half-reaction but NOT the comparative value for a
net reaction.
 If the half-reaction is used in reverse (for RA or oxidation potential)
then the sign on the Electrical potential should also be changed.
(similar to Hess’s law)
 Multiplying a half-reaction by a coefficient in order to balance has
NO effect on the Electrical potential. (do not multiply the potential –
different than Hess’s law)
When adding the half-reactions to get the net reaction, the voltages are also
added to find the net voltage. OR Eonet(cell) = Eocathode - Eoanode
 A positive cell potential implies that the reaction is spontaneous.
 A negative cell potential implies that it is non-spontaneous.
Eg.
The values in the data booklet are for conditions listed on the bottom of the
page! Changing concentration may affect the potential but only slightly
since amount of chemical does not affect cell potential. Only if the
concentrations are very weak would we see a drastic change.
Should the chemicals be in a different state, it will affect the Eo values and
the data booklet ones should NOT be used (Electrolysis uses).
Corrosion and Cathodic Protection:
 You will describe the methods and devices used to prevent corrosion; i.e.,
physical coatings and cathodic protection
What is rust? http://en.wikipedia.org/wiki/Rust
Very often we use metal or metal components that can readily come into
contact with water from the environment, soil, and air. This poses the
problem of corrosion. Metal corrosion costs companies a lot of money in
replacement costs.
There are numerous ways to help slow down or prevent corrosion:
1. Protective coatings on the metals. Eg. Plastic wraps, galvanizing.
2. Sacrificial anodes. A more reactive metal (stronger RA – anode) is
used on the surface of the metal structure being protected(or in circuit
with the metal structure). This reactive metal is called a sacrificial
anode because it is more likely to react (stronger RA) before the other
metal therefore protecting the metal structure. As the sacrificial anode
reacts it releases electrons that supply current to the metal structure
and force it to be the cathode. This helps suppress the reaction of the
metal structure.
The sacrificial anode has to be maintained and replaced frequently
(depending on the amount of exposure and corrosion).
Eg. Motor boat motors, hull of ships, wrap on piping, anode beds for
piping.
3. Impressed current. If the negative terminal of a battery or power
supply is connected to the metal being protected, and the positive
terminal is connected to an inert electrode, then a current flows
through the metal forcing it to be the place of reduction (cathode)
rather than oxidation when electrolyte (ground water) is present.
It requires a constant supply of current.
Eg. pipelines
ELECTROLYTIC CELLS –
 You will recognize that predicted reactions do not always occur; e.g., the
production of chlorine gas from the electrolysis of brine
-a cell that uses electrical energy to create chemical energy (connected to a
power supply)
-Still uses SOA/SRA when predicting the reaction however, the reactants
can be influenced by the amount of voltage applied to the cell.
-Eo should be a negative value which indicates the minimum amount of
voltage required to cause the reaction to occur.
Eg. Potassium iodide
*Watch out for the chloride anomaly in electrolytic cells. Pg645
Secondary Cells:
-the anode and cathode switch sides since the reactions reverse. The anode
is now positive and the cathode is negative. Anode is connected to
positive post of power supply and cathode to the negative so electrons are
supplied for the reduction reaction.
USES OF ELECTROLYSIS:
READ PG 646-650 and make notes of the uses (*chlor-alkali)
1. Production of Pure Elements
2.
Refining of metals
3.
Electroplating
Do #16,18 pg649
Do #3,4,7,9,10,11,12,14,15 on pg 651
Electrolytic Cell Stoich:
 You will calculate mass, amounts, current and time in single voltaic and
electrolytic cells by applying Faraday’s law and stoichiometry.
Oxidation-reduction equations are quantitative/stoichiometric and
therefore mathematical manipulations can be done. For cells, this involves
the amount of current and time that the power is supplied.
Quantity
Symbol
Unit
Equivalent units
Electrical Potential
V, Eo
Volt
Joule/coulomb
(voltage)
Charge
q
Coulomb
-Current
I
Ampere
Coulomb/second
*see pg 2 of data booklet
*page 3 is faraday’s constant: 1 mol of electrons = 9.65 x 104C of charge
1 electron = 1.6 x 10-19 C
New Formula****
ne = It/F
(because ne = q/F and I = q/t or q=It so putting those together gives you the above
formula)
Faraday’s Law of Electrolysis;
1. Mass of an element produced at an electrode is directly proportional
to the electric charge transferred in the cell. Q=It
2. Mass of an element produced at an electrode depends on the # of
electrons transferred per entity and the molar mass of the element.
Eg. 1. In a copper refinery, pure metal is produced by electrolysis. If a cell operates for
40.0 min at 12.0A, how much copper is deposited at the cathode
b) If the same operation was applied to lead, how much lead would be deposited?
c) If the same operation was applied to aluminum, how much aluminum would
be deposited?
2.
In an electrolytic cell 21.3 g of Cl2 was produced from molten NaCl in 80
minutes. Find the average current used. Which electrode would this be deposited
at?
3.
How long must a 0.500 A current run to plate 5.87 g of Ni from molten NiCl2(l).
At what electrode does this happen?
b) How much chlorine gas is formed at the other electrode by the same process?