Chemistry Chapter 19 oxidation-reduction reactions

Oxidation-Reduction Reactions
Chapter 19

Oxidation and Reduction
Section 1

Oxidation-reduction (redox) reactions involve transfer of electrons
Oxidation – loss of electrons
Reduction – gain of electrons
Both half-reactions must happen at the same time
Can be identified through understanding of oxidation numbers

Oxidation States
Oxidation number assigned to element in molecule based on distribution of electrons in molecule
There are set rules for assigning oxidation numbers

Chromium gives great example of different oxidation numbers
Different oxidation states of chromium have different colors
Chromium (II) chloride = blue
Chromium (III) chloride = green
Potassium chromate = yellow
Potassium dichromate = orange

Oxidation
Oxidation reactions in which the atoms or ions of an element experience an increase in oxidation state
Ex. combustion of metallic sodium in atmosphere of chlorine gas

Sodium ions and chloride ions made during exothermic reaction form cubic crystal lattice
Sodium cations are ionically bonded to chloride anions

Formation of sodium ions shows oxidation b/c each sodium atom loses an electron to become sodium ion
Oxidation state represented by putting oxidation number above symbol of atom and ion

Oxidation state of sodium changed from 0 (elemental state) to +1 (state of the ion)
A species whose oxidation number increases is oxidized
Sodium atom oxidized to sodium ion

Reduction
Reduction reactions in which the oxidation state of an element decreases
Ex. Chlorine in reaction with sodium
Each chlorine atom accepts e- and becomes chloride ion
Oxidation state decreases from 0 to -1

A species that undergoes a decrease in oxidation state is reduced
The chlorine atom is reduced to the chloride ion

Oxidation and Reduction as a Process
Electrons are made in oxidation and acquired in reduction
For oxidation to happen during chemical reaction, reduction must happen as well
Number of electrons made in oxidation must equal number of electrons acquired in reduction
Conservation of mass

Transfer of e- causes changes in oxidation states of one or more elements
Oxidation-reduction reaction any chemical process in which elements undergo changes in oxidation number
Ex. When copper oxidized and NO3- from nitric acid is reduced

Part of the reaction involving oxidation or reduction alone can be written as a half-reaction
Overall equation is sum of two half-reactions
Number of e- same of oxidation and reduction, they cancel and don’t appear in overall equation

Electrons lost in oxidation appear on product side of oxidation half-reaction
Electrons gained in reduction appear as reactants in reduction half-reaction

When copper reacts in nitric acid 3 copper atoms are oxidized to Cu+2 ions as two nitrogen atoms are reduced from a +5 oxidation state to a +2 oxidation state

If no atoms in reaction change oxidation state, it is NOT a redox reaction
Ex. Sulfur dioxide gas dissolves in water to form acidic solution of sulfurous acid

When solution of NaCl is added to solution of AgNO3, an ion-exchange reaction occurs and white AgCl precipitates

Redox Reactions and Covalent Bonds
Substances with covalent bonds also undergo redox reactions
Unlike ionic charge, oxidation number has no physical meaning
Oxidation number based on electronegativity relative to other atoms to which it is bonded in given molecule
NOT based on charge

Ex. Ionic charge of -1 results from complete gain of one electron by atom
An oxidation state of -1 means increase in attraction for a bonding electron
Change in oxidation number does not require change in actual charge

When hydrogen burns in chlorine a covalent bond forms from sharing of two e-
Two bonding e- in hydrogen chloride not shared equally
The pair of e- is more strongly attracted to chlorine atom because of higher electronegativity

As specified by Rule #3, chlorine in HCl is assigned oxidation number of -1
Oxidation number for chlorine atoms changes from 0
So chlorine atoms are reduced

From Rule #1, oxidation number of each hydrogen atom in hydrogen molecule is 0
By Rule #6, oxidation state of hydrogen atom in HCl is +1
Hydrogen atom oxidized

No electrons totally lost or gained
Hydrogen has donated a share of its bonding electron to chlorine
It has NOT completely transferred that electron
Assignment of oxidation numbers allows determination of partial transfer of e- in compounds that are not ionic
Increases/decreases in oxidation number can be seen in terms of completely OR partial loss or gain of e-

Reactants and products in redox reactions are not limited to monatomic ions and uncombined elements
Elements in molecular compounds or polyatomic ions can also be redoxed if they have more than one non-zero oxidation state
Example: copper and nitric acid

Nitrate ion, NO3-, is converted to nitrogen monoxide, NO
Nitrogen is reduced in this reaction
Instead of saying nitrogen atom is reduced, we say nitrate ion is reduced to nitrogen monoxide

Balancing Redox Equations
Section 2

Equations for simple redox reactions can be balanced by looking at them
Most redox equations require more systematic methods
Equation-balancing process needs use of oxidation numbers
Both charge and mass are conserved
Half-reactions balanced separately then combined

Half-Reaction Method
Also called ion-electron method
Made of seven steps
Oxidation numbers assigned to all atoms and polyatomic ions to determine which species are part of redox process
Half-reactions balanced separately for mass and charge
Then added together

Sulfur changes oxidation state from -2 to +6
Nitrogen changes from +5 to +4
Other substances deleted

In this example, sulfur is being oxidized

To balance oxygen, H2O must be added to left side
This gives 10 extra hydrogen atoms on that side
So, 10 H atoms added to right side
In basic solution, OH- ions and water can be used to balance atoms

Electrons added to side having greater positive net charge
Left side has no net charge
Right side has +8
Add 8 electrons to product side
(oxidation of sulfur from -2 to +6 involves loss of 8 e-)

Nitrogen reduced from +5 to +4

H2O added to product side to balance oxygen atoms
2 hydrogen ions added to reactant side to balance H atoms

Electrons added to side having greater positive net charge
Left side has net charge of +1
1 e- added to this side balancing the charge

This ratio is already in lowest terms
If not, need to reduce
Multiply oxidation half-reaction by 1
Multiple reduction half-reaction by 8
Electrons lost = electrons gained

Each side has 10H+, 8e-, and 4H2O
They cancel

The NO3- ion appeared as nitric acid in original equation
Only 6 H ions to pair with 8 nitrate ions
So, 2 H ions must be added to complete this formula
If 2 H ions added to left side, then 2 H ions must be added to the right side

Sulfate ion appeared as sulfuric acid in original equation
H ions added to right side used to complete formula for sulfuric acid

Sample Problem

The iron (II), iron (III), manganese (II), and 2 H ions in original equation are paired with sulfate ions
Iron (II) sulfate requires 10 sulfate ions
Sulfuric acid requires 8 sulfate ions
To balance equation, 18 sulfate ions must be added to each side

On product side, 15 of these form iron (III) sulfate, and 2 form manganese (II) sulfate
Leaves 1 sulfate unaccounted for
Permanganate ion requires the addition of 2 potassium ions to each side
These 2 K ions form potassium sulfate on product side

Oxidizing and Reducing Agents
Section 3

Reducing agent substance that has the potential to cause another substance to be reduced
They love electrons
Attain a positive oxidation state during redox reaction
Reducing agent is oxidized substance

Oxidizing agent substance that has the potential to cause another substance to be oxidized
Gain electrons
Attain a more negative oxidation state during redox reactions
Oxidizing agent is reduced substance

Strength of Oxidizing and Reducing Agents
Different substances compared and rated on relative potential as reducing/oxidizing agents
Ex. Activity series – related to each element’s tendency to lose electrons
Elements lose electrons to positively charged ions of any element below them in series

The more active the element the greater its tendency to lose electrons
Better a reducing agent it is
Greater distance between two elements in list means more likely that a redox reaction will happen between them

Fluorine atom most highly EN atom
Is also most active oxidizing agent
b/c of strong attraction for its own e-, fluoride ion is weakest reducing agent
Negative ion of strong oxidizing agent is weak reducing agent

Positive ion of strong reducing agent is weak oxidizing agent
Ex. Li
Strong reducing agents b/c Li is very active metal
When Li atoms oxidize they produce Li+ ions
Li+ ions unlikely to reacquire e-, so it’s weak oxidizing agent

Left column of each pair also shows relative abilities of metals listed to displace other metals
Zinc, ex., is above copper so is more active reducing agent
Displaces copper ions from solutions of copper compounds
Copper ion is more active oxidizing agent than Zn

Nonmetals and others are included in series
Any reducing agent is oxidized by oxidizing agents below it
Ex. F2 displaces Cl-, Br-, and I- from their solutions

Autooxidation
Some substances can be both reduced and oxidized
Ex. Peroxide ions – O2-2 has relatively unstable covalent bond

Each O atom has oxidation number of -1
Structure represents intermediate oxidation state between O2 and O2-2
So, peroxide ion is highly reactive

Hydrogen peroxide, H2O2, contains peroxide ion
Decomposes into water and oxygen as follows

Hydrogen peroxide is both oxidized AND reduced
Oxygen atoms that become part of gaseous oxygen molecules are oxidized (-1  0)
Oxygen atoms that become part of water are reduced (-1  -2)

Autooxidation  a process in which a substance acts as both an oxidizing agent and a reducing agent
The substance is self-oxidizing and self-reducing

Bombardier Beetle
Defends itself by spraying its enemies with an unpleasant hot chemical mixture
Catalyzed autooxidation of H2O2 produces hot oxygen gas
Gas gives insect ability to eject irritating chemical from abdomen

Electrochemistry
Section 4

Redox reactions involve energy changes
b/c these involve electron transfer, the net release or net absorption of energy can happen in the form of electrical energy instead of heat
This allows for many practical applications
Also makes possible predictions and comparisons of oxidizing and reducing abilities of different substances
Electrochemistry branch of chemistry that deals with electricity-related applications of redox reactions

Electrochemical Cells
Redox reactions involve transfer of electrons from oxidized to reduced substance
If two substances in contact, a transfer of energy as heat happens as well

Zinc strip in contact with copper (II) sulfate solution
Zinc strip loses e- to copper (II) ions in solution
Copper (II) accept electrons and fall out of solution as copper atoms
As electrons are transferred, energy released as heat (rise in temperature)

If we separate substances the e- transfer comes with transfer of electricity instead of heat
One way to separate is with porous barrier
Prevents metal atoms of one half-reaction from mixing with metal atoms of the other

Ions in two solutions can move through porous barrier
e- can be transferred from one side to the other through connecting wire
Current moves in circuit so movement of e- balanced by movement of ions in solution

Altering system in Figure 19-7 so electrical current is produced involves separating copper and zinc (Figure 19-8)
Zn strip is in aqueous solution of ZnSO4
Cu strip is in solution of CuSO4

Both solutions are electrolytes
Electrode conductor used to establish electrical contact with nonmetallic part of a circuit, such as an electrolyte
Zn and Cu strips are electrodes

Half-cell single electrode immersed in a solution of its ions
Zn strip in ZnSO4 (aq) is anode  where oxidation takes place
Cu strip in CuSO4 (aq) is cathode  where reduction takes place

Copper half-cell written as Cu2+/Cu
Zinc half-cell written as Zn2+/Zn
Two half-cells together make electrochemical cell system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produces chemical change

Electrochemical cell can be represented by following notation:
Anode І Cathode
Cell made of zinc and copper would be written Zn І Cu
2 types of electrochemical cells
Voltaic (also called galvanic)
electrolytic

Voltaic Cells
Voltaic cell  if redox reaction in electrochemical cell happens spontaneously and produces electrical energy
Cations in solution reduced when they gain e- at surface of cathode to become metal atoms

Half reaction for cathode:
Half reaction for the anode:

e- given up at anode pass along external wire to cathode

Movement of e- through wire must be balanced by movement of ions in solution
Anions move to anode to replace negative e- moving away

Cations move to cathode as positive charge lost through reduction

Dry cells are common sources of electrical energy
They are voltaic cells
Three most common types:
Zinc-carbon battery
Alkaline battery
Mercury battery

Zinc-Carbon Dry Cells
Consist of zinc container (anode)
Filled with moist paste of MnO2, graphite, and NH4Cl
When external circuit closed, Zn oxidized at negative electrode (anode)

e- move across circuit and reenter cell through carbon rod (cathode)
MnO2 reduced in presence of water

Alkaline Batteries
Do not have carbon rod cathode
Allows them to be smaller
Uses paste of Zn metal and KOH instead of solid metal anode
Half-reaction at anode:

Reduction half-reaction (at cathode) same as Zn-C dry cell

Mercury Batteries
Tiny batteries in hearing aids, calculators, etc.
Anode half-reaction same as alkaline dry cell

Cathode half-reaction is as follows:

Electrolytic Cells
Some redox reactions do not happen spontaneously but drive by electrical energy
Electrolysis process in which electric current used to produce a redox reaction
Electrolytic cell  electrochemical cell where electrical energy is required to produce a redox reaction and bring about a chemical change

Electrode of cell connected to negative terminal of battery gets an excess of e- and becomes the cathode of electrolytic cell
Electrode of cell connected to positive terminal of battery loses e- to battery and becomes the anode
e- pump simultaneously supplying e- to cathode and recovering from anode
This energy input from battery drives electrode reactions in electrolytic cell

Voltaic cell has copper cathode and zinc anode
If battery connected to positive terminal contacts copper electrode and negative terminal contacts zinc electrode, e- flow in opposite directions

Battery forces cell to reverse its reaction
Zn becomes cathode
Cu becomes anode

Half-reaction at anode (Cu oxidized):
Half-reaction at cathode (Zn reduced):

Important Differences Between Voltaic and Electrolytic Cells
The anode and cathode of an electrolytic cell are connected to a battery or other direct-current source, whereas a voltaic cell serves as a source of electrical energy.
Electrolytic cells have electrical energy from external sources creating nonspontaneousredox reactions. Voltaic cells have spontaneousredox reactions occurring.
Electrolytic cell have electrical energy converted to chemical energy. Voltaic cells have chemical energy converted to electrical energy

Electroplating
Metals like Cu, Ag, Au are difficult to oxidize
In electrolytic cell, these metals form ions at anode that are easily reduced at a cathode
This allows solid metal from one electrode to be deposited on the other electrode
Electroplating an electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface

Electroplating cell has solution of a salt of plating metal
Has object to be plated (cathode)
Has piece of plating metal (anode)

Silver-plating cell has solution of soluble silver salt and silver anode
Cathode is object to be plated
Silver anode connected to positive electrode of battery or other source of direct current
Object to be plated connected to negative electrode

Silver ions reduced at cathode according to following equation
Silver ions deposited as metallic silver when electrons flow through circuit

Metallic silver is removed from anode as ions
Silver atoms oxidized at anode as follows
This maintains Ag+ concentration of solution

Rechargeable Cells
Rechargeable cell combines redox chemistry of both voltaic and electrolytic cells
When cell converts chemical energy to electrical energy it is a voltaic cell
When recharged, it operates as an electrolytic cell converting electrical to chemical energy

The standard 12V car battery is set of 6 rechargeable cells
Anode in each cell is lead submerged in solution of H2SO4

Anode half-reaction is as follows

e- move through circuit to cathode where PbO2 is reduced as follows

Next redox reaction for discharge cycle is as follows

A car’s battery produces electric energy needed to start engine
Sulfuric acid (present as ions) is consumed
Lead (II) sulfate builds up as white powder on electrodes
Once car is running, half-reactions are reversed by voltage produced by alternator
Pb, PbO2 and H2SO4 are regenerated
Battery can be recharged as long as all reactants are present and all reactions are reversible

Electrode Potential
There are 2 electrodes, Zn and Cu
These 2 metals each have different tendencies for accepting electrons
Reduction potential tendency for half-reaction of a metal to happen as a reduction half-reaction in an electrochemical cell

There are 2 half-cells
Strip of Zn in solution of ZnSO4
Strip of Cu in CuSO4
Electrode potential difference in potential between an electrode and its solution

When these 2 half-cells are connected and reaction begins, difference in potential is observed between the electrodes
This voltage is a measure of energy required to move certain electric charge between electrodes
Potential difference measured in volts

Voltmeter connected across the Cu І Zn voltaic cell measures potential difference of about 1.10 V when solution concentrations are each 1 M

Potential difference measured across voltaic cell roughly equals sum of electrode potentials for the 2 half-reactions
Easy to measure voltage across voltaic cell
No way to measure individual electrode potential directly
b/c there is no transfer of e- unless both anode and cathode are connected in complete circuit

Relative value for potential of half-reaction can be determined by connecting it to standard half-cell as reference
This is called a standard hydrogen electrode, or SHE

Platinum electrode dipped in 1.00 M acid solution
Surrounded by hydrogen gas at 1 atm pressure and 25° C
Other electrodes ranked according to ability to reduce H under these conditions

Anodic reaction for SHE is described by forward half-reaction in following equation
Cathodic half-reaction is reverse
Random potential of 0.00 V assigned to both of these reactions
Any voltage measurement is credited to the half-cell connected to SHE
Standard electrode potential, E⁰ a half-cell potential measured relative to a potential of zero for the SHE

Electrode potentials expressed as potentials for reduction
Provide reliable indication of tendency of a substance to be reduced
Positive E⁰values indicate H more willing to give up e- than other electrode
Half-reactions with positive reduction potentials are favored

Half-reactions with negative reduction potentials are not favored
These half-reactions prefer oxidation over reduction
Negative E⁰ values indicate that the metal or other electrode is more willing to give up e- than hydrogen

When half-reaction is written as oxidation rxn, sign of electrode potential changed as shown for redox half-rxns for Zn

To measure reduction potential of Zn half-cell, it is connected to SHE
Potential difference is -0.76 V
Negative number indicates e- flow through external circuit from Zn electrode (Zn oxidized) to H electrode (H ions reduced)

Copper half-cell paired with SHE gives E⁰ of +0.34 V
This positive number indicates Cu2+(aq) more readily reduced than H+(aq)

Standard electrode potentials can be used to predict if redox reactions will happen naturally
A naturally occurring rxn will have positive value for E⁰cell according to the following equation

When evaluating the 2 half-rxns of a cell, the half-rxn with more negative E⁰ is the anode
Oxidation happens at the anode, so half-cell rxn is reverse of reduction rxn in Table 19-4
When rxn reversed, actual half-cell potential is negative of E⁰
The total potential of a cell is calculated by subtracting standard reduction potential for rxn at anode (E⁰anode) from standard reduction potential for rxn at cathode (E⁰cathode)

Example
Cells made from Fe in solution of Fe(NO3)3 and Ag in solution of AgNO3
Table 19-4 gives

Fe in Fe(NO3)3 is anode b/c it has lower reduction potential
Ag in AgNO3 is cathode
Overall cell reaction is

Reduction of Ag ions multiplied by 3 so number of e- lost in that half-reaction equals number of e- gained in oxidation of iron
Standard reduction potentials for anode and cathode are as follows
NOTE!
When a half-reaction is multiplied by a constant, the E⁰ value is NOT multiplied.

So…
Potential for this cell can be calculated as follows
If calculated value for E⁰cell were negative, the reaction would NOT happen naturally
It would not happen in a voltaic cell
It could be made to happen in an electrolytic cell

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Chemistry Chapter 19 oxidation-reduction reactions

by Renee Ariesen on Jul 06, 2010

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Chemistry Chapter 19 oxidation-reduction reactions — Presentation Transcript