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Unit 3 - Electrochemistry > Chapter 13 - Redox Reactions > Flashcards

Chapter 13 - Redox Reactions Flashcards

1
Q

What is the basis of an electrochemical reaction?

A

Electron transfer.

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2
Q

What type of reactions are the most common in both living and nonliving systems?

A

Electrochemical reactions.

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3
Q

Some examples of electrochemical processes are:

A

Photosynthesis, cellular respiration, and metabolism.

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4
Q

Do gold and silver naturally exist as a pure element?

A

Yes.

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5
Q

What are ores?

A

Rocks in which most elements exist mixed with other elements.

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6
Q

What is metallurgy?

A

extracting metals from their naturally occurring compounds.

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7
Q

What is reduction (historical)?

A

Producing metals from their compounds.

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8
Q

What is oxidation (historical)?

A

The process of a substance reacting with oxygen (like corrosion or combustion).

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9
Q

What does a reducing agent do?

A

Causes or promotes reduction.

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10
Q

What does an oxidizing agent do?

A

Causes or promotes oxidation.

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11
Q

What are the most common reducing agents used in metallurgical processes?

A

Carbon monoxide, carbon, and hydrogen.

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12
Q

what is the main difference between corrosion and combustion?

A

Corrosion occurs more slowly.

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13
Q

Corrosion could be considered the opposite of…

A

Metallurgy.

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14
Q

What is a half reaction?

A

A balanced chemical equation that shows either the loss or gain of electrons of a substance.

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15
Q

The gain of electrons is:

A

reduction.

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16
Q

The loss of electrons is:

A

oxidation.

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17
Q

The total number of electrons lost must be equal to

A

the total number of electrons gained.

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18
Q

Are oxidation and reduction separate processes?

A

Yes.

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19
Q

Do oxidation and reduction occur simultaneously?

A

Yes.

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20
Q

Oxidation-reduction reactions are regularly called:

A

Redox reactions.

21
Q

Steps for creating a half reaction in an acidic solution:

A
  1. Write the reactants and products.
  2. Balance everything other than water and hydrogen.
  3. Add water to balance the oxygen atoms.
  4. Because its acidic, you have hydrogen atoms available to balance the hydrogen.
  5. Add electrons.
22
Q

Tug-of-war analogy and spontaneity:

A

Each entity pulls on the same electrons and tries to win by pulling the hardest. If one entity wins, a spontaneous reaction occurs.

23
Q

A reducing agent will undergo

A

oxidation.

24
Q

An oxidizing agent will undergo

A

reduction.

25
Q

When did the OA and RA terms originate?

A

During early metallurgy.

26
Q

What was known as the only oxidizing agent early on>

A

Oxygen.

27
Q

Oxidation and reduction are ______ while OA’s and RA’s are ______.

A

Processes; substances.

28
Q

Where is the strongest oxidizing agent on a redox table?

A

Top left.

29
Q

Where is the strongest reducing agent in a redox table?

A

Bottom right.

30
Q

What does the redox spontaneity rule state?

A

A spontaneous reaction only occurs if the OA is above the RA in a redox table.

31
Q

If its an aqueous solution add

A

water.

32
Q

If its an acidic solution you can add

A

hydrogen ions.

33
Q

How to predict redox reactions in solution (five step method):

A
  1. List all entities present and classify them as possible OA’s or RA’s. Don’t label spectator ions.
  2. Choose the SOA and write its equation (reduction).
  3. Choose the SRA and write its equation (oxidation).
  4. Balance the electrons then put the equations together to get your net ionic equation.
  5. See if its spontaneous using the spontaneity rule.
34
Q

If the table does not have the half-reaction equation, what should you do?

A

Create one yourself (by doing the thing where you add water, etc.)

35
Q

What is an oxidation state?

A

The apparent net electric charge that an atom would have if electron pairs in covalent bonds belonged entirely to the more electronegative atom.

36
Q

What is an oxidation number?

A

A positive or negative number corresponding to the oxidation state assigned to an atom in a covalently bonded entity. Just shows the electrons that can be gained/ given up.

37
Q

All atoms in elements oxidation numbers are:

A

Zero.

38
Q

Hydrogen in all compounds except hydrides oxidation number:

A

Positive one.

39
Q

Disproportionation.

A

Where a substance is both oxidized and reduced at the same time.

40
Q

Oxidation number of oxygen in compounds except peroxides:

A

Negative two.

41
Q

In H2O2, oxygen’s oxidation number is:

A

Negative one.

42
Q

What is hydrogen’s oxidation number when it acts as an anion?

A

Negative one.

43
Q

Halogen’s oxidation number is usually:

A

Negative one.

44
Q

An increase in oxidation number indicates:

A

Oxidation.

45
Q

A decrease in oxidation number indicates:

A

Reduction.

46
Q

Balancing using oxidation numbers steps:

A
  1. Assign oxidation numbers to all of the atoms and ions and note the ones that change.
  2. See how many electrons have been travelled. For example, a change from -2 to +4 is 6 e-.
  3. Multiply by the subscript. For example, in O2, you would multiply the electrons by two. Within an atom, just look at the thing that changed. For example, in H2S and the sulfur changed, keep the electrons as is.
  4. Multiply the electrons of each of the things that changed so that they are equal. This is because the electrons gained and lost should be the same.
  5. The numbers you multiplied become the coefficients. for the reactants.
  6. Balance everything else based on these numbers.
  7. At the end you may need to add water molecules and hydrogen ions if its in an acidic solution.
47
Q

When is the equivalence point reached during titration?

A

When the exact number of moles of reactants have been used up.

48
Q

When is endpoint reached during titration?

A

When the next drop of excess reactant is added noted by a colour change.

49
Q

Redox stoichiometry:

A

Find the SOA and SRA and write those reactions out. Then, convert your solution from mol/L to mols. Then do the mol to mol ratio.

Unit 3 - Electrochemistry (2 decks)

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