REDOX

Redox - Concepts to Master• Give everyday examples of REDOX

reactions.• Why must redox reactions contain both

oxidation and reduction half reactions?• What happens to electrons in oxidation?• What happens to electrons in reduction?• What is the oxidation state of an element in

its pure form?• How do assign oxidation numbers?• Write a sentence that contains the words:

oxidation, electrons, and cation.• Write a sentence that contains the words:

reduction, electrons, and anion.• What is the difference between oxidation

and an oxidizing agent?• What is the difference between reduction

and a reducing agent?• Be able to write half reactions given a

chemical equation.• Be able to determine the net ionic equation.• Be able to pick out the spectator ions.

• What are strong electrolytes?• Give examples of strong electrolytes?• Be able to balance a redox equation.• What two things are conserved during a redox

reaction?• Be able to use Table J to determine spontaneous

reactions.• How can you tell if a reaction is spontaneous?• Why is a salt bridge required in a voltaic cell?• What are the physical differences between voltaic

and electrolytic cells?• Where does oxidation take place in a voltaic and

electrolytic cell?• Where does reduction take place in a voltaic and

electrolytic cell?• Which electrode loss mass in electrolysis?• Are voltaic or electrolytic cells spontaneous?• Chemical energy is changed to electrical energy

in which type of cell?• Electrical energy is changed to chemical energy

in which type of cell?• In which direction do electrons flow in the voltaic

cell?• In which direction do electrons flow in the

electrolytic cell?

Vocab• Anions• Anode• Cathode• Cations• Electrochemical cell• Electrochemistry• Electrolysis• Electrolytes• Electrolytic cell• Electroplating• Galvanic cell• Half reactions• Negative terminal

• Oxidation• Oxidation number• Oxidation state• Oxidizing agent• Positive terminal• Principle species• Reducing agent• Reduction• Salt bridge• Spectator ions• Spontaneous• Standard• Voltaic cell

Labs• Oxidation of Mn• Alchemy

Electrochemistry and Redox• The branch of chemistry that studies the relationships

between electric forces and chemical reactions.• Electrochemistry plays an important part in our everyday

lives. – Responsible for the rusting of iron– Allows for the purification many metals – Allows for the plating common metals with silver or

gold– Explains how batteries power iPods

• In order to understand electrochemistry, we must first examine Reduction/Oxidation reactions (also known as REDOX reactions).

REDOX• Type of chemical reaction• A redox (reduction-

oxidation) reaction is one in which one or more atoms change oxidation number from reactants to products.

– The number of electrons lost by one atom must equal the number of electrons gained by another atom.

– Therefore both reduction and oxidation must occur in a redox reaction.

Ion Types - Cations

• Sodium has 3 PELs• The third PEL has 1 electron. • To be more stable Na gives away 1 electron.

(This is easier than receiving 7 electrons.) • Sodium will lose 1 electron which means that it

now has 10 electrons and 11 protons.

• It has 1 less electron than proton so it’s charge is +1.

• An electron has been lost so OXIDATION has occurred.

• Na --> Na+1 + 1e• The above reaction is an

oxidation half reaction.Cats are +

• Oxygen has 2 PELs.• The second PEL has 6 electrons. • To be more stable O will receive 2 more

electrons from another atom. (This is easier than giving away all 6 electrons.)

• Oxygen will receive 2 more electrons which means it now has 10 electrons and 8 protons.

• It has 2 more electrons than protons so it’s charge is -2.

• Electrons have been gained so reduction has occurred.

• O + 2e --> O-2

• The above reaction is a reduction half reaction.

Ion Types - Anions

+ 2 e

Assigning Oxidation Numbers

• What is the oxidation number of sulfur in magnesium sulfate?

• The formula for magnesium sulfate is MgSO4

• The oxidation number of SO4 is -2 • The oxidation number of O4 = 4 x (-2) = -8• Since the total oxidation number of the sulfate

ion is -2, the oxidation number of sulfur is: -8 + S = -2

• S = +6

Assigning Ox #

• Pure elements or molecules, oxidation # = 0

Na(s) H2(g) O2(g) N2(g) F2(g) Hg(s)

• What is the oxidation number on N in NO2?• N+4

• What is the oxidation number of S in SO4-2?

• S+6

Oxidation has occurred when…

• Electrons have been lost

• An atom has become a positive ion (cation)

• An atom’s oxidation number has increased

2 Na(s) + Cl2(g) → 2 NaCl (s)

Na0 becomes Na+1 so it can combine with Cl.

Note that this reaction is also a synthesis reaction.

Reduction has occurred when…

• Electrons have been gained

• An atom has become a negative ion (anion)

• An atom’s oxidation number has decreased

2 Na(s) + Cl2(g) → 2 NaCl (s)

Cl0 becomes Cl-1 so it can combine with Na.

Half Reactions• The part of the reaction that only

involves oxidation OR reduction.

• Oxidation half reaction:

Na0 → Na+1 + 1e

• Reduction half reaction:

Cl20 + 2e → 2Cl -1

Fe(s) + CuSO4(aq) → FeSO4 (aq) + Cu(s)

In the reaction….

• Iron is oxidized by giving up electrons to form FeSO4 .– How many?– Write the half reaction– Name FeSO4

• Copper is reduced by receiving the electrons to form Cu(s). – How many?– Write the half reaction Note this is also a Note this is also a

single replacement reactionsingle replacement reaction

Oxidation half reaction

Fe(s) Fe2+(aq) + 2e

Reduction half reaction

Cu2+(aq) + 2e Cu(s)

Overall

Fe(s) + Cu2+(aq) + 2e Fe2+

(aq) + Cu(s) + 2e

All electrons lost in oxidation half-reaction must be gained in a reduction half-reaction.

LEO the lion says GER

• Losing Electrons is Oxidation

• Gaining Electrons is Reduction

OIL RIG

• Oxidation Is Losing

• Reduction Is Gaining

2Mg(s) + O2(g) → 2MgO (s)

H2O

In the reaction….

• Magnesium is oxidized by giving up electrons to form MgO .– How many?– Write the half reaction

• Oxygen is reduced by receiving the electrons to form MgO. – How many?– Write the half reaction

Synthesis ReactionSynthesis Reaction

The overall net charge on both sides of the equation is zero.

Oxidation half reaction

Mg(s) Mg2+(s) + 2e

Reduction half reaction

O2(g) + 2e O-2(s)

Overall

Mg(s) + O2(g) + 2e Mg2+(s) + O-2

(s) + 2e

All electrons lost in oxidation half-reaction must be gained in a reduction half-reaction.

Label each reaction as Oxidation or Reduction

1. Co0 → Co+2 + 2 e

2. N0 + 3 e → N-3

3. Br0 + 1 e → Br-1

4. Ni0 → Ni+3 + 3 e

5. Ag0 → Ag+1 + 1 e

6. Sr0 → Sr+2 + 2 e

7. Se0 + 2 e → Se-2

8. Sb0 + 3 e → Sb-3

9. Fe0 → Fe+2 + 2 e

10.W0 → W+6 + 6 e

Do metals or nonmetals tend to gain e and be reduced?

Do metals or nonmetals tend to lose e and be oxidized?

(secret) Agents

• Oxidizing Agent (oxidizer)– The substance that causes the oxidation of other

substances.– Location where reduction occurred– It itself is reduced.

• Reducing Agent– The substance that causes the reduction of other

substances.– Location where oxidation occurred– It itself is oxidized.

Label the element in the reactions as an oxidizing agent or a reducing agent:

1. Co0 → Co+2 + 2 e

2. N0 + 3 e → N-3

3. Br0 + 1 e → Br-1

4. Ni0 → Ni+3 + 3 e

5. Ag0 → Ag+1 + 1 e

6. Sr0 → Sr+2 + 2 e

7. Se0 + 2 e → Se-2

8. Sb0 + 3 e → Sb-3

9. Fe0 → Fe+2 + 2 e

10.W0 → W+6 + 6 e

Do metals tend to be oxidizers or reducing agents?

Do non-metals tend to be oxidizers or reducing agents?

Net ionic equations• The reason to write a chemical equation is to express what

is happening in a chemical reaction.

• In net ionic equations:– the focus is on the ions involved in a reaction--the

principal species– ignoring those spectator ions that really don't get

involved.

• HCl (aq) + NaOH (aq) → NaCl (aq) + H2O

• HCl, NaOH, and NaCl are all strong electrolytes. As such, they dissociate (or separate) completely into their ions when dissolved in water (“in solution”).– HCl is H+ and Cl– – NaOH is Na+ and OH–

– NaCl is Na+ and Cl–

Net ionic equations continued

• HCl (aq) + NaOH (aq) → NaCl (aq) + H2O could be

rewritten…

• H+(aq) + Cl–(aq) + Na+

(aq) + OH–(aq) →Na+

(aq) + Cl–(aq) + H2O(l)

• Notice that Na+ and Cl– never really react. They are

floating around at the beginning and still floating around at

the end. They are just “spectators”.

• So an equation that represents what is really going on

would be:

• H+(aq) + OH–

(aq) → H2O(l)

Rules for writing net ionic equations

• You MUST start with an equation that includes the physical state: – (s) for solid, – (l) for liquid, – (g) for gas, and – (aq) for aqueous solution– Table F

• Only break up the (aq) substances. • Only break up strong electrolytes. • Delete any ions that appear on both sides of

the equation.

Strong electrolytes

• Ionic compounds (SALTS)• Strong Acids

– HCl– HBr– HI

– HNO3

– HClO3

– HClO4

– H2SO4

• Strong Bases– NaOH– KOH– LiOH

– Ba(OH)2

– Ca(OH)2

Net Ionic Equations Practice

– Write an equation for the reaction of a solution of BaCl2 with a solution of Na2SO4 (predict products)

– Assign state of matter (consult Table F)

– BaCl2(aq) + Na2SO4(aq) → NaCl(aq) + BaSO4(s)

– Since barium sulfate is insoluble, we can NOT separate it into its ionic components.

– 2Na+1(aq) + SO4

-2(aq) + Ba+2

(aq) + 2Cl-1(aq) -----> 2Na+1

(aq) + 2Cl-1(aq) + BaSO4(s)

– Na and Cl are on both sides of the equation so they are the spectator ions and can be deleted.

– SO4-2

(aq) + Ba+2(aq) ----> BaSO4(s)

Identify the Spectator Ions in these ionic equations

• Al(s) + 6H1+(aq) + 6Cl1-

(aq) ----> Al3+(aq) + 6Cl1-

(aq) + 3H2(g)

– 6Cl1-(aq)

• K(s) + H2O(l) ----> K1+(aq) + OH1-

(aq) + H2(g)

– none

• 2Ag(s) + Cu2+(aq) + SO4

2-(aq) ----> 2Ag1+

(aq) + Cu(s) + SO42-

(aq)

– SO42-

(aq)

1. Solid calcium + aqueous zinc nitrate

2. Solid iron + aqueous tin sulfate 

3. Aqueous Iron (III) iodide + chlorine gas

4. Aqueous sodium carbonate + aqueous zinc chloride

Write the net ionic equations for each and list the spectator ions.

Balancing Redox Reactions

1. The unbalanced net ionic equation is above.2. Separate into half reactions.

Cu → Cu+2

Ag+1 → Ag3. Add electrons appropriately to the oxidation and reduction

reactions.Cu → Cu+2 + 2eAg+1 + 1e→ Ag

4. Equalize the electrons in the half reactions.4. Cu → Cu+2 + 2e5. (Ag+1 + 1e→ Ag)x2 2Ag+1 + 2e→ 2Ag

5. Add the half reactions together. Cu → Cu+2 + 2e2Ag+1 + 2e→ 2Ag

Balance

2Ag+1 + Cu → Cu+2 + 2Ag

6. Check that atom # and charges are equal on both sides

2Ag+1 + Cu → Cu+2 + 2Ag

2+0=2 2+0=22 Ag 2 Ag 1 Cu1 Cu

Balance

Table J – Activity Series (again?)

• Used to determine if a single replacement reaction will occur.– Each element from the list displaces from a compound any of the

elements below it.

• Elements will oxidize elements below them and reduce elements above them. The higher elements are…

• Stronger reducing agents• Lose electrons more easily• More easily oxidized

• Also used to determine if a reaction occurs spontaneously.– If the elementelement appears higher in the table than the ionion, then the

reaction will occur spontaneously.

Electron Transfer = Redox Reactions

• The electron transfer seen in a redox reaction can generate electricity which can then do “work”.

• HOW?– Electrochemical Cell

• A system that contains two electrodes separated by an electrolyte (salt bridge).

• Galvanic cell or voltaic cell• Electrolytic cell

Galvanic Cell or Voltaic Cell (your common battery)

• A copper rod (called an electrode) is immersed in a solution of copper (II) sulfate (called an electrolyte) and connect by a wire to a zinc rod immersed in a solution of zinc sulfate.

• Looking at the Activity Chart (table J), we can see that zinc is the stronger reducing agent (it is higher on the list) therefore, the zinc will be more easily oxidized (lose electrons) and the copper will be reduced (gain electrons).

• In order for reaction to occur, electrons must flow from the zinc rod to the copper rod (as indicated by the arrows). In this example, the zinc rod is called the anode and the copper rod is called the cathode.

An Ox sat on a Red Cat• The cathode is the electrode where

reduction occurs. – This should be the element that is lower on

Table J.

• The anode is the electrode where oxidation occurs.– This should be the element that is higher on

Table J.

• As the reaction proceeds, solid zinc metal will become aqueous zinc ions and aqueous copper ions will become solid copper metal. This causes…– the zinc sulfate solution to become positively charged – the copper sulfate to become negatively charged

• In order to maintain neutral charges in both electrolytes, they are connected by a salt bridge. – The salt bridge is simply a solution of spectator ions that can

move into each electrolyte through a porous membrane. – The salt bridge allows the migration of ions between the half

cells and therefore keeps the solutions neutral.

Galvanic Cell / Voltaic Cell & the Salt Bridge

• Eventually, the Zn metal is gone and the battery dies.

• The anode loses mass.

• The cathode gains mass.

Galvanic Cell or Voltaic Cell

Electrolytic Cell• If electricity is used to

force electrons to move in the non-spontaneous direction (forcing the stronger oxidizing agent to act as a reducing agent and vice versa) the cell is called an electrolytic cell.

• In the electrolytic cell, the zinc rod is now the cathode (where the reduction occurs) and the copper rod is now the anode (where the oxidation occurs), opposite of the galvanic cell.

• This is how we recharge batteries.

• Due to the battery, electrons are forced through the cell in the opposite direction that they want to go. The electrons will continue to flow in the reverse direction for as long as electricity to supplied.

Electrolytic Cell

Galvanic / Voltaic A galvanic cell is one in which a redox reaction takes place spontaneously to produce electricity. Chemical energy is changed to electrical energy.  

Electrolytic CellAn electrolytic cell is one in which electrolysis occurs, a compound is decomposed by passing electricity through it. The reaction occurring is not spontaneous, it is forced by applying a voltage (electricity). Electrical energy is converted to chemical energy

      

Applications of Electrolysis (Electrolytic Cells)

Galvanic/Voltaic Cells Electrolytic Cells

chemical energy    electrical energy electrical energy     chemical energy

two half-cells with separate electrolytes and a salt bridge (or porous barrier).

electrodes usually in the same electrolyte

chemical reaction is spontaneous chemical reaction is forced by applying a voltage - it is not spontaneous

anode - negative terminal cathode - positive terminal

anode - positive electrode cathode - negative electrode

Electrons are released into the wire from the anode so the anode is the negative terminal of the battery

electrons flow from the negative battery terminal (or the electrical outlet) to the cathode making it the negative terminal

electrons flow from the negative terminal to the positive terminal (anode to cathode)

electrons flow from the positive terminal to the negative terminal (anode to cathode)

uses - batteries uses - extract Al from Al2O3 (extraction

of metals from ore), electroplating, purifying metals

Voltaic

Electrolytic