Download Hebden V.2 – Oxidation Numbers

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Hebden V.1 – Introduction
Basics:
Electrochemistry: the branch of chemistry which is
concerned with the conversion of chemical energy to
electrical energy and vice versa
Oxidation: a half reaction in which a species loses
electrons – LEO= loss of electrons - oxidation
Reduction: a half reaction in which a species gains
electrons – GER = gain of electrons - reduction
Hebden V.1 – Introduction
Basics:
Half-reaction: the reaction that occurs in one half cell of an
electrochemical cell
Electrochemical cell: a system which produces
electricity – uses chemical reactions, involves the loss and
gaining of electrons
Hebden V.1 – Introduction
Oxidizing agents
 are reduced –
 oxidizing agents gain electrons –
 oxidizing agents cause reduction reactions
 half reaction has electrons on the reactant side –
Reducing agents
 are oxidized –
 reducing agents loose electrons –
 reducing agents cause oxidizing reactions
 half reaction has electrons on the product side
Hebden V.1 – Introduction
Recognizing oxidation and reduction
Oxidation = in the products - oxidation number increases,
number of oxygen atoms increases
Reduction = in the products - oxidation number decreases,
number of oxygen atoms decreases
Hebden V.2 – Oxidation Numbers
Oxidation numbers: charge an atom would have if the species
containing the atom where made up of ions
The sum of all positive charges and negative charges must equal the
overall charge on the species
Step 1: write the formula for the molecule
Step 2: write the known oxidation numbers or charges below each
species in the molecule – put in an x for the unknowns
Step 3: write a math equation showing the sum of all the species as
equal to the overall charge of the ion
Hebden V.2 – Oxidation Numbers
1.
2.
3.
4.
5.
6.
oxygen is virtually always –2 (in peroxides it is –1)
the alkali metals are usually +1
the alkali earth metals are usually +2
the halogens are usually –1 (Cl, Br, I, F)
Polyatomic ions have an overall charge that will be shown like OHNeutral molecules do not have a charge shown – it is zero – H4P2O7 has a
charge of 0
7. All atoms have charge of 0
8. Hydrogen in all compounds (except hydrides like LiH) is +1 – look to see if H
is at the beginning or end of the compound
9. The oxidation number for a monatomic ion is the same as the charge Cl- =
oxidation number -1
Hebden V.2 – Oxidation Numbers
Example
Find the oxidation # for P in H4P2O7
H4 P2 O7
Charge on one atom
+1 x -2
Charge on all atoms
+4 2x -14
0 (overall charge on
molecule)
Make an equation
+4 + 2x + (-14) = 0 solve to find x
Hebden V.3 – Predicting Spontaneity of a Redox Reaction
On the table
1. oxidizing agents are on the right
2. reducing agents are on the left
3. oxidizing agents are strongest at the top
4. reducing agents are strongest at the bottom
5. some elements may be oxidizing or reducing agents
6. use double arrows to show the reactions can go forward or
reverse when you are talking about an isolated half reaction but
when you are using a half reaction as part of a redox reaction
use a one way arrow to show if it is reduced or oxidized
Hebden V.3 – Predicting Spontaneity of a Redox Reaction
Reactions are spontaneous only if
1. the reactant to be reduced (the oxidizing agent) is above the
reactant to be oxidized (the reducing agent).
2. both reactants are on the same side = no reaction
3. the reducing agent is above the oxidizing agent = no reaction
Check for H+ - for many reactions – H must be present for the
reaction to happen – check it
Reactions happen only if there is a reactant to be reduced (an
oxidizing agent) which is above the reactant to be oxidized (a
reducing agent).
Hebden V.3 – Predicting Spontaneity of a Redox Reaction
Tables of relative strengths
To make tables to show the relative strengths of oxidizing or
reducing agents, or the relative tendencies of oxidizing or reducing
agents:
a) Identify the strongest oxidizing agent (SOA) and the
strongest reducing agent (SRA) – check the table
b) Write half reactions with the metal ions on the right and the
metal atoms on the left
c) Use spontaneity to determine which are OA and which are RA
– if the reaction is spontaneous, the OA is above the RA on
the table
Hebden V.4 - Balancing Half-Reactions
1.Write the chemical formulas for the reactants and
products
2. Balance all the atoms except O and H
3. balance the O’s by adding H2O’s(l)
4. balance the H by adding H+(aq)
5. balance the charge on each side by adding e- and
cancel anything that is on both sides
Hebden V.4 - Balancing Half-Reactions
For basic solutions only
6. add OH-(aq)to both sides equal in number to the H+
present
7. combine H+(aq) and OH-(aq) on the same side to form
H2O(l)
8. cancel equal amounts of H2O(l) and anything else
that is the same on both sides
Hebden V.5 – Balancing Redox Equations Using HalfReactions
You don’t have to worry about oxidation numbers
Rules:
1. Separate the reactants and products into two groups
2. For each group (each ½ reaction) balance the “O’s” using
H2 O
3. balance the “H’s” using H+
4. Balance the charge using e5. balance the redox reaction by conserving electrons
6. Add the two halves together
7. If necessary convert to OH- (basic solutions)
Hebden V6 – Balancing Redox Equations Using Oxidation
Numbers
Balancing Redox Reactions using Oxidation Numbers
1. Identify those species that change oxidation numbers
2. Make sure electrons gained = electrons lost (balance
oxidation number charge) – this gives the coefficients for
the backbone of the reaction
3. Add waters to balance the “O’s”
4. Add H+ to balance the “H’s”
5. If you need to change to basic conditions add OH- to
each side and reduce or cancel the waters (basic
solutions)
Hebden V6 – Balancing Redox Equations Using Oxidation
Numbers
1.
2.
3.
4.
5.
6.
List all the entities in the reaction
Classify them as RA, OA, or both
Find the SOA and SRA – use the table
Choose the SOA and write the reduction half reaction
Choose the SRA and write the oxidation half reaction
Balance the number of electrons lost and gained in the half reaction
equations by multiplying one or both the equations by the coefficient of
the number of electrons in the opposite equation
7. If the reaction is acidic, add H+ to balance the hydrogen atoms
8. Add the two balanced half reaction equations to get a net ionic equation
9. Use the spontaneity rule and the table to decide if the reaction is spontaneous or not.
Hebden V.6 - Redox Titrations
For titrations you need something that will change colour
when it is oxidized – for an oxidizing agent use MnO4- for a reducing agent use ITo find the concentration of an element that is easily
oxidized use KMnO4
To find the concentration of an element that is easily
reduced NaI or KI
The process for titration questions is essentially the same as
for acid – base problems
Hebden V.7 - Electrochemical cells
Electrode – conductor at which a half-reaction occurs
Anode – conductor at which oxidation occurs
- receives electrons from substance being oxidized
- electrode towards which anions travel
Cathode – conductor at which reduction occurs
- gives electrons away to a substance being
reduced
- electrode towards which cations travel
Hebden V.7 - Electrochemical cells
To make an electrochemical cell you need:
Parts: 2 beakers with solutions
u-tube
wire
2 electrodes – metals
Hebden V.7 - Electrochemical cells
Electrons flow
1. from the anode to the cathode
2. through the wire
3. only ions flow in the solutions
4. the number of electrons in the reduction reaction
must equal the number of electrons in the oxidation
reaction
Hebden V.9 - Standard Reduction Potentials
Electric potential = voltage = work done per electron transferred
1. Write the oxidation and reduction half reactions with their
potentials (from the table)
2. E0cell = E0RED - E0ox
3. If you have to reverse the way the reaction is written to show
the reduction or oxidation you must reverse the sign on the E0
4. If E0cell is positive = spontaneous
5. If E0cell is negative = non-spontaneous
1.
2.
3.
4.
Hebden V.10 - Selecting Preferred Reactions
identify all the ions – separate any ionic compounds
start on the reduction side of the table – go to the
upper left – run your finger down until you find the first
match with a species on your list This is your oxidizing
agent
start on the oxidizing side of the table – go to the lower
right – run your finger up until you find the first match
with a species on your list This is your reducing agent
Write the two half reactions and add them together
Hebden V.11 - Applied Electrochemistry
No notes Read this section, make sure you understand
and can relate some examples.
Hebden V.12 - Corrosion of Metals: Cause and
Prevention
The cathode region has lots of oxygen
The anode region is oxygen poor
To prevent corrosion:
1. paint
2. use a cathode
Cathodic protection = the process of protecting a substance from
unwanted oxidation by connecting it to a substance having a higher
tendency to oxidize.
When choosing a cathode pick something with a higher tendency to
oxidize. Zn works with many metals
Hebden V.13 - Electrolysis
Electrolysis is the process of supplying electrical energy to a
molten ionic compound or a solution containing ions so as to
produce a chemical change
Electrolytic cell – apparatus in which electrolysis occurs
Hebden V.13 - Electrolysis
To make and electrolytic cell
Parts – container
Solution with ions
External power source
2 electrodes
Hebden V.13 - Electrolysis
The reactions that happen are not spontaneous so the
reduction reaction (the OA) will be below the oxidation
reaction (the RA) on the table
The electrodes must be inert – Pt or C work well
Anode reaction – oxidized
Cathode reaction – reduced
Hebden V.13 - Electrolysis
Choosing preferred reactions
1. list all the entities present
2. Determine if the conditions are acidic or neutral –
always assume neutral unless one of the chemicals is an
acid
3. find the highest reduction reaction (OA)
4. find the lowest oxidation reaction (RA)
5. Write the half reactions and add them together
6. Add the electric potentials to find out how much voltage
must be in the external power source
Hebden V.13 - Electrolysis
Electroplating
The cathode is made out the material that will receive the metal
plating
The electroplating solution contains ions of the metal to be plated
onto the cathode
The anode may be made of the same metal which is to be plated onto
the cathode but an inert electrode may be used.
Electrons go from the anode to the cathode so hook up the power
supply so the cathode will get the electrons.
Make your anode a metal
Make sure your ions are in solution with something that will not
make it precipitate NO3 works well