Unit 5: Oxidation and Reduction

Question 1

What was oxidation and reduction initially referred to?

Answer 1

Oxidation - An element combining with oxygen to produce an oxide.

Reduction - A substance losing oxygen

Question 2

What is oxidation and reduction?

Answer 2

Oxidation - Losing electrons

Reduction - Gaining electrons

Question 3

What are oxidizing agents and reducing agents?

Answer 3

Oxidizing agents - gain electrons (take them)

Reducing agents - lose electrons (give them)

Question 4

What are net and total ionic equations?

Answer 4

Total - Break apart aqueous substances into ions

Net - Remove spectator ions.

Question 5

Redox of moleculars compounds

Answer 5

Molecular compounds are not ionic, they share electrons. This unequal sharing determines what is oxidized and reduced (use the electronegativity values).

Ex: H2 + O2 -> H2O, Oxygen is more electronegative that hydrogen, so the oxygen holds electrons more often, and is reduced (partially gains electrons).

Question 6

What is a half reaction?

Answer 6

Balanced to have an equal number of electrons.

Ex. S + 2e- -> S^-2

Question 7

How do you tell whether a reaction is spontaneous or not?

Answer 7

If the oxidizing agent is higher than the reducing agent on a table of reduction half reactions, then it is spontaneous. And vice versa.

Question 8

What the the word corrosion from?

Answer 8

Latin: to gnaw

Question 9

When does corrosion usually occur?

Answer 9

When water is in contact with the metal in the presence of oxygen.

Question 10

What can corrosion cause?

Answer 10

Things to become stuck.
Holes in metal.
Poor conductors.
Lost magnetic ability.
Weaker metals.

Question 11

How can corrosion be prevented? How?

Answer 11

Painting/coating: prevents water and oxygen from contacting metal. Oil and grease can help as well.
Galvanizing: coat in zinc, prevents contact and sacrificial metal.
Sacrificial metal: More reactive metal is placed in the vicinity of the metal that needs corrosion protection. Does not necessarily need to coat metal.
Alloys: Can have better properties, such as not rusting, or using a s as ridiculous metal (like nickel and chromium).
Cathodic protection: sacrificial metal connected via electric current that transfers electrons.

Question 12

What is an alloy?

Answer 12

A metal combined with another substance to create new properties.

Question 13

Oxidation number of a free uncombined element?

Answer 13

0

Question 14

Oxidation number of oxygen combined with something else?

Answer 14

-2

Question 15

Total of all oxidation numbers in a compound?

Answer 15

0

Question 16

Oxidation number of monatomic ion?

Answer 16

Equal to charge, ex: Fe(III) = +3

Question 17

Oxidation number of H combined with something else?

Answer 17

+1, aside from metal hydrides, in which case -1.

Question 18

How do you balance with oxidation numbers?

Answer 18

Assign numbers
Identify what is oxidized and reduced
Write change in oxidation number by square brackets
Multiply to make changes equal
Put coefficients in front
Check overall balance and make changes

Question 19

What reactions normally are redox, aren’t, or sometimes are?

Answer 19

Usually: combustion and single displacement
Not: neutralization and double displacement
Sometimes: synthesis and decomposition

Question 20

What are the characteristics of redox reactions?

Answer 20

Either oxygen atoms, hydrogen atoms, or electrons are transferred.
Both oxidation and reduction must be present.
Reciprocity: If you took the reverse reaction of oxidation, the result would be reduction, and vice versa.

Question 21

What is patina?

Answer 21

A layer of oxide formed on the surface of metal

Question 22

Why doesn’t copper rust away like iron?

Answer 22

Iron flakes and peels when it rusts, exposing more iron.
Copper leaves a layer of oxide on the surface that protects the rest.

Question 23

Oil rig acronym.

Answer 23

Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

Question 24

What are the six types of chemical reactions?

Answer 24

Synthesis, single displacement, double displacement, decomposition, combustion, neutralization.

Question 25

What chemical reactions are sometimes redox and sometimes not?

Answer 25

Synthesis and decomposition.

Question 26

What chemical reactions are usually redox?

Answer 26

Combustion and single displacement.

Question 27

What chemical reactions are not considered redox?

Answer 27

Neutralization and double displacement.

Question 28

Steps to balance with half reactions.

Answer 28

1. Write the separate half reactions for the reduction and oxidization (include the electrons)
2. Balance the atoms in the half reactions.
   Use H2O and H+ to balance in acidic solutions. Use H2O and OH- to balance in basic solutions.
3. Multiply each half reaction by an appropiate coefficient to make the electron changes equal.
4. Add the half reactions and subtract terms that appear on both sides of the equation.

Question 29

What is an easy trick to balance a half reaction with H2O and OH-?

Answer 29

For each O a side should have, add 2 OH-. Have half as much H2O as OH-.

Question 30

Who found out that different metal strips (electrodes) can generate an electrical current in a conducting solution?

Answer 30

Leigh Galvani.

Question 31

Who developed the first electrochemical cells?

Answer 31

Alessandro Volta. He first used copper and zinc in a salt solution.

Question 32

Recall the beaker examples. What does each beaker contain?

Answer 32

A half reaction.

Question 33

Using the beaker example, how does an electrochemical cell work?

Answer 33

Two beakers contain half reactions (one anode, one cathode). They are connected by a piece of wire and a salt bridge (wire conducts electrons, electrolytic solution replenishes electrons and extends lifetime). Wire and salt bridge create a circuit that allow the flow of electrons.

Anode (metal part) produces positive ions, and contains electrons. These electrons flow to the cathode, which has a positive charge.

The salt bridge (internal circuit) releases negative ions to the anode solution to keep the solution electrically neutral.

Question 34

What is an external circuit?

Answer 34

Where the electrical work is done. Ex, a tv remote powered by batteries is an external circuit.

Question 35

What is an internal circuit?

Answer 35

At the anode, ions are produced and go into the solution. This causes a build-up of positive ions in this solution, which can hinder the reactions. The salt bridge (internal circuit) keeps the solution electrically neutral.

Question 36

In what direction do electrons flow?

Answer 36

From the anode (negative) to the cathode (positive).

Question 37

How do you calculate cell voltage?

Answer 37

Locate the two half cells reactions in the standard reduction potential table. The reaction that is higher on the table will reduce. The other will oxidize (write the given reaction in reverse). Balance for the number of electrons exchanged, but do not change voltages.
Add the voltages together (invert the sign of the oxidizing reaction). Remember to add the two half reactions together.

Question 38

What is an electrolytic cell?

Answer 38

Electrochemical cells are spontaneous, and produce electricity. An electrolytic cell flows in the reverse direction (not spontaneous) so electricity must be added (think phone batteries).

Question 39

What are some uses of electrolytic cells?

Answer 39

Rechargeable batteries.
Electroplating.
Purifying metal.

Question 40

How does metal purifying work with electrolytic cells?

Answer 40

Electrons travel from the anode to the cathode, reducing the metal ions into metal atoms. The metal plates onto the cathode as it forms. The non-metal ions (which are negative) migrate towards the anode and release electrons as they oxidize to form atoms.

Ex. Sodium chloride can become solid sodium and chlorine gas.

Question 41

Where does reduction and oxidation occur in an electrochemical cell?

Answer 41

Reduction at cathode
Oxidation at anode