Chapter 17 Electrochemistry

Question 1

a redox reaction

Answer 1

One that entails changes in oxidation number (or oxidation state) for one or more of the elements involved.

Question 2

the oxidation number of an neutral atom in an element is

+ex

Answer 2

Zero
ex: Cl2

Question 3

Oxidation number

*ex with +/- on

Answer 3

The total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. The iron ion Fe3+, for example, has an oxidation number of +3 because it can acquire three electrons to form a chemical bond, while the oxygen ion O2− has an oxidation number of −2 because it can donate two electrons.

Question 4

In an electronically neutral substance, the sum of the oxidation numbers is

Answer 4

0

Question 5

oxygen ON is

Answer 5

-2 but -1 in peroxide (ANY PROXIDE)

Question 6

Consequential to these rules, the sum of oxidation numbers for all atoms in a molecule is equal to

Answer 6

the charge on the molecule

Question 7

What are the oxidation numbers?

Answer 7

the oxidation number for nitrogen is +5 and that for oxygen is −2, summing to equal the 1− charge on the molecule:What a

Question 8

Oxidation

Answer 8

Lose electron

Question 9

Reduction

Answer 9

Gain in electron

Question 10

balancing equations for aqueous redox reactions steps:

Answer 10

1. Write skeletal equations for the oxidation and reduction half-reactions.
2. Balance each half-reaction for all elements except H and O
3. Balance each half-reaction for O by adding H2O.
4. Balance each half-reaction for H by adding H+.
5. Balance each half-reaction for charge by adding electrons.
6. If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
7. Add the two half-reactions and simplify.
8. If the reaction takes place in a basic medium, add OH− ions the equation obtained in step 7 to neutralize the H+ ions (add in equal numbers to both sides of the equation) and simplify

Question 11

galvanic cells (or voltaic cells) (5)

+ what is included in the setup of the cell

Answer 11

• those in which a spontaneous redox reaction takes place
• comprised of two half-cells, each containing the redox conjugate pair (one solid one aq ex: solid cu/cu2+aq and solid Ag/Ag2+)
• An external circuit is connected to each half-cell at its solid foil, meaning the Cu and Ag foil each function as an electrode.
• An external circuit is connected to each half-cell at its solid foil, meaning the Cu and Ag foil each function as an electrode.
• To keep the reactants separate while maintaining charge-balance, the two half-cell solutions are connected by a tube filled with inert electrolyte solution called a salt bridge

Question 12

anode

What occurs here and which process?

Answer 12

the electrode at which oxidation occurs
positive end, where electrons leave

Question 13

cathode

Answer 13

The electrode where reduction occurs
negative terminal or electrode through which electrons enter

Question 14

The redox reactions in a galvanic cell occur only at

Answer 14

the interface between each half-cell’s reaction mixture and its electrode.

Question 15

In the galvanic cell, positive charge build up at

Answer 15

the annode

Question 16

In the galvanic cell, negative charge build up at

Answer 16

the cathode

Question 17

Why a salt bridge is needed

Answer 17

Without the salt bridge, the solution in the anode compartment would become positively charged and the solution in the cathode compartment would become negatively charged,because of the charge imbalance. The attractive and repulsive forces will prohibit the flow of electrons within the cell.

Question 18

Difference between Galvanic and electrolytic cells

Answer 18

Galvanic cells derives its energy from spontaneous redox reactions, while electrolytic cells involve non-spontaneous reactions and thus require an external electron source like a DC battery or an AC power source.

Question 19

cell notations (3)

Answer 19

• The relevant components of each half-cell are represented by their chemical formulas or element symbols
• All interfaces between component phases are represented by vertical parallel lines; if two or more components are present in the same phase, their formulas are separated by commas
• By convention, the schematic begins with the anode and proceeds left-to-right identifying phases and
  interfaces encountered within the cell, ending with the cathode

Question 20

inert electrode (2)

Why we use this+how does it work

Answer 20

• This is required when neither member of the half-cell’s redox couple can reasonably function as an electrode, which must be electrically conductive and in a phase separate from the half-cell solution.
• simply provide or accept electrons to redox species in solution

Question 21

active electrodes

Answer 21

Electrodes constructed from a member of the redox couple

Question 22

When measured for purposes of electrochemistry, a potential reflects the

Answer 22

driving force for a specific type of charge transfer process, namely, the transfer of electrons between redox reactants.

Question 23

cell potentials, Ecell (2)

What it is+formula

Answer 23

the difference in potential between two half-cells

Question 24

standard cell potential, E°cell

Answer 24

a cell potential measured when both half-cells are under standard-state conditions (1 M concentrations, 1 bar pressures, 298 K)

Question 25

standard hydrogen electrode (SHE) (3)

What it its signifigance/what is its conditions and compoents/formula

Answer 25

One particular half-cell that serve as a universal reference for cell potential measurements, assigning it a potential of exactly 0 V.
- A typical SHE contains an inert platinum electrode immersed in precisely 1 M aqueous H+ and a stream of bubbling H2 gas at 1 bar pressure, all maintained at a temperature of 298 K

Question 26

standard electrode potential, E°X (EX) for a half-cell X is defined as

Answer 26

the potential measured for a cell comprised of X acting as cathode and the SHE acting as anode all under standard conditions

Question 27

Good oxidizing agents have —- standard reduction potentials

Answer 27

high

Question 28

good reducing agents have — standard reduction potentials.

Answer 28

low

Question 29

In a reduction standard potential table, the highest values (most negative) represent the—–. This means that the more negative values are more likely to be the —-

Standard reduction potential measures the tendency for a given chemical species to be reduced.

Answer 29

stronger reducing agents.
anode

Question 30

an increased value of E° corresponds to an increased driving force behind

Answer 30

the reduction of the species (hence increased effectiveness of its action as an oxidizing agent on some other species)

Question 31

Negative values for electrode potentials are simply a consequence of assigning a value of 0 V to the SHE, indicating the reactant of the half-reaction is

Answer 31

a weaker oxidant than aqueous hydrogen ions.

Question 32

E°cell is positive when E°cathode — E°anode, and so any redox reaction in which the oxidant’s entry is above the reductant’s entry is predicted to be —–.

Answer 32

> spontaneous

Question 33

The spontaneous oxidation of copper by lead(II) ions is not observed,and so the reverse reaction, the oxidation of lead by copper(II) ions, is predicted to occur —

Answer 33

spontaneously

Question 34

reversing the direction of a redox reaction effectively interchanges the — and the cell potential is now calculated…..

Answer 34

identities of the cathode and anode half-reactions, and so the cell potential is calculated from electrode potentials in the reverse subtraction order than that for the forward reaction

Question 35

In the case of a redox reaction taking place within a galvanic cell under standard state conditions, essentially all the work is associated

Answer 35

transferring the electrons from reductant-to-oxidan

Question 36

The work associated with transferring electrons is determined by

Answer 36

where n is the number of moles of electrons transferred, F is Faraday’s constant, and E°cell is the standard cell potential.

Question 37

E° and K relating formula (2)

+what does this equation tell us?

Answer 37

This equation indicates redox reactions with large (positive) standard cell potentials will proceed far towards completion, reaching equilibrium when the majority of reactant has been converted to product

Question 38

Nernst equation

What is it and when do we use it?

Answer 38

This equation describes how the potential of a redox system (such as a galvanic cell) varies from its standard state value

Question 39

A convenient form of the Nernst equation for most work is one in which values for the fundamental constants (R and F) and standard temperature (298) K), along with a factor converting from natural to base-10 logarithms:

Answer 39