Reaction Types

Question 1

Define and give a general reaction for:

a Brønsted-Lowry acid

Answer 1

A Brønsted-Lowry acid is any species capable of donating a proton to the solution and results in an increase in hydronium ion concentration, hence a decrease in pH.

Note: This is the definition that the AP Chemistry test uses for any acid in general.
Example acid reaction: HA + H₂O ⇒ A^- + H₃O⁺

Question 2

Define and give a general reaction for:

a Brønsted-Lowry base

Answer 2

A Brønsted-Lowry base is any species capable of accepting a proton from the solution and results in an increase in hydroxide ion concentration, hence an increase in pH.

Note: This is the definition that the AP Chemistry exam uses for any general base.
Example base reaction: NH₃ + H₂O ⇒ NH₄⁺ + OH^-

Question 3

Is CH₃COOH an acid or a base?

Answer 3

CH₃COOH is an acid. It loses a proton to solution.

Acid Equation: CH₃COOH + H₂O ⇒ CH₃COO^- + H₃O⁺

Question 4

Is NH₃ an acid or a base?

Answer 4

NH₃ is a base. It accepts a proton (or creates an OH-) in solution.

Base Reaction: NH₃ + H₂O ⇒ NH₄⁺ + OH^-

Question 5

Define and give an example of:

a strong acid

Answer 5

A strong acid is one which dissociates completely in solution.

For a monoprotic acid (one proton per molecule), each mole of acid in solution results in one mole of protons in solution as well.
Ex: HCl is a classic strong acid (monoprotic).

Question 6

Define and give an example of:

a strong base

Answer 6

A strong base is one which dissociates completely in solution.

For a monobasic compound (one hydroxide per molecule), each mole of base in solution results in one mole of hydroxide ions in solution as well.
Ex: NaOH is a classic strong base (monobasic).

Question 7

Place these in order of increasing acid strength:

H₂O; NH₃; HF; CH₄

Answer 7

CH₄ < NH₃ < H₂O < HF

Acidity, in general, is a measure of how easily that substance will donate a proton into solution.

Question 8

Why is this the ranking of increasing acid strength?

CH₄ < NH₃ < H₂O < HF

Answer 8

CH₄ < NH₃ < H₂O < HF

These molecules are an increasing series going from left to right across the second row of the Periodic Table. Traveling left to right across the Periodic Table, acidity always increases.

The reason for this is polarity. Chemically, these acids are: a hydrogen atom bound to a central atom. As the central atom becomes more electronegative, the bonds with hydrogen become more polar. More polar bonds are easier to dissociate in aqueous solution. Hence, the more electronegative the central atom, the more easily it donates protons, and the more acidic it is.

Question 9

Place these in order of increasing acid strength:

HCl; HF; HI; HBr

Answer 9

HF < HCl < HBr < HI

Recall: acid strength is determined by how easily the substance will donate a proton into solution.

Question 10

Why is this the order of increasing acid strength?

HF < HCl < HBr < HI

Answer 10

HF < HCl < HBr < HI

These molecules are a series going down a column of the Periodic Table. As you travel down a column in the Periodic Table, acidity increases.

The reason for this is atomic size. Larger atoms can carry negative charges more easily, so the I^- ion is more stable than the F^- ion. The more stable the conjugate base, the stronger the acid.

Question 11

Please list six common strong acids.

Answer 11

1. HI (Hydrogen Iodide)
2. HBr (Hydrogen bromide)
3. HCl (Hydrogen chloride)
4. HNO₃ (Nitric Acid)
5. HClO₄ (Perchloric Acid)
6. H₂SO₄ (Sulfuric Acid)

While there are several other acids to choose from, these are the ones most commonly used in most chemistry courses, including on the AP Chem exam.

Question 12

Please list seven common strong bases.

Answer 12

1. NaOH (sodium hydroxide)
2. KOH (potassium hydroxide)
3. NH₂^- (amide ion)
4. H^- (hydride ion)
5. Ca(OH)₂ (calcium hydroxide)
6. Na₂O (sodium oxide)
7. CaO (calcium oxide)

While there are several other bases to choose from, these are the ones most commonly used in most chemistry courses, including on the AP Chem exam.

Question 13

Define and give an example of:

a polyprotic acid

Answer 13

A polyprotic acid can donate more than one proton to a solution.

Ex: H₂SO₄ (can donate 2 protons)

Question 14

Define and give an example of:

an amphoteric substance

Answer 14

An amphoteric substance can act as an acid or a base depending on the solution.

The classic example is water.
Acting as a base: H₂0 + HA ⇔ A^- + H₃0⁺
Acting as an acid: H₂0 + B- ⇔ BH + OH^-

Question 15

Define:

conjugate acid-base pairs

Answer 15

Conjugate acid-base pairs are molecules which differ via the presence or absence of a proton. The protonated form of the molecule is the conjugate acid, the deprotonated form is the conjugate base (Brønsted-Lowry acid definition)

Generic equation: HA + H₂O → A^- + H₃O⁺
HA is an acid and A^- is its conjugate base. Similarly, H₂O is acting as a base, with H₃O⁺ acting as its conjugate acid.

Question 16

What is the conjugate base of acetic acid, CH₃COOH?

Answer 16

The acetate ion, CH₃COO^-

An acid’s conjugate base is the deprotonated remainder of the molecule’s acid reaction.
The general acid neutralization reaction is HA + OH^- ⇒ A^- + H₂O
where HA is the acid, and A- is the conjugate base.

Question 17

What is the product of the reaction of a strong acid and a strong base?

Answer 17

Salt and water, according to the general equation

HA + BOH ⇒ AB + H₂0

It is possible to add acid and base and NOT create water (Lewis acid/base pairs) but there will always be a salt formed. Additionally, when you use Bronsted-Lowry as the acid/base definition (like the AP Chem exam does), it is safe to assume that water is always formed as well.

Question 18

What is the purpose of a titration experiment?

Answer 18

The purpose of a titration experiment is to discover the concentration of an unknown acid or base, by neutralizing it with a measured quantity of a base or acid solution whose concentration is known.

In titrations, the known solution is the titrant and the unknown solution is the analyte.

Question 19

Define:

the equivalence point

Answer 19

The equivalence point is the point at which every molecule of acid has been neutralized by a molecule of base.

[H⁺ ions]_original = [OH^- ions]_added
In a titration, this is where all of the unknown (analyte) has been completely neutralized by the known (titrant), giving a neutral pH.

Question 20

If a weak acid is titrated with a strong base, what is the pH at the equivalence point?

Ex: CH₃COOH + NaOH ⇒
NaCH₃COO + H₂O

Answer 20

More than 7, slightly basic.

The equivalence point is a point at which the amount of equivalents of acid and equivalents of base from the analyte and titrant are equal.
In the example: a weak base (acetate ion, CH₃COO^-) is being formed, hence the pH at the equivalence point will be above 7.

Question 21

Define and give an example of:

an indicator

(as it relates to titrations)

Answer 21

In a titration, the indicator changes the color of the solution to indicate that the equivalence point pH has been reached.

Ex: Phenolphthalein goes from colorless to fuchsia between pH 8.3-10
Ex: Methyl red goes from red to green to yellow between pH 4.4-5.2-6.2

Question 22

1. What are the requirements for an indicator?
2. Name several common ones.

Answer 22

1. An indicator is a molecule that must change color visibly in a set pH range, usually a range of about 2 pH units. To select an indicator for a specific titration, find one whose pK_a is roughly equal to the pH of the titration’s equivalence point.
2. Some common indicators include:
   
   • Methyl red (pH range 4.4-6.2)
   • Thymol blue (8.0-9.6)
   • Azolitmin (4.5-8.3)

Question 23

Explain and give the pH value for:

The equivalence point in a strong acid/strong base titration

Answer 23

The equivalence point is a point at which the amount of equivalents of acid and equivalents of base from the analyte and titrant are equal. For a strong acid/strong base combination, this happens at a pH of 7. At the equivalence point:

V_A * N_A = V_B * N_B

Where
V_A = volume of acid in L
N_A = normality of acid in #equivalents*mol/L
V_B = volume of base in L
N_B = normality of base in #equivalents*mol/L

Question 24

Define and give the common use for:

an analyte

Answer 24

An analyte is an acid or base whose concentration is determined by a titration.

A titrant (known concentration) is used to bring the unknown (analyte) solution up to a standard pH (usually neutral 7). In this way, the concentration of the analyte can be calculated.

Question 25

Define and give an example of the use for:

a titrant

Answer 25

A titrant is a strong acid or base with a known concentration.

In a titration experiment the titrant is used to neutralize, or bring to a known pH, an unknown (analyte) solution. In this way, the concentration of the unknown can be calculated.

Question 26

In the below titration of the acid H₂A, at which point does
[H₂A] = [HA^-]?

Answer 26

Point A

Point A is the first half-equivalence point, where one-half as many equivalents of base have been added as acid molecules that were in the solution to begin with, so one half of the acid molecules have been neutralized, and [H₂A] = [HA^-].

Question 27

In the below titration of the acid H₂A, what is the earliest point where the solution is entirely A^2-?

Answer 27

Point D

A^2- production is not complete until 2 full equivalents of base have been added. This must be at the second equivalence point, which is Point D.

Question 28

What is the chemical process happening to the Ag atom in this reaction?

Ag(s) ⇒ Ag⁺(aq) + e^-

Answer 28

The Ag atom is undergoing oxidation.

The mnemonic to use in redox (reduction-oxidation) reactions is OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

In this case, the Ag atom is losing an electron, so it is being oxidized.

Question 29

What is the chemical process happening to the Fe³⁺ ion in this reaction?

Fe³⁺(aq) + e^- ⇒ Fe²⁺(aq)

Answer 29

The Fe³⁺ ion is undergoing reduction.

The mnemonic to use in redox (reduction-oxidation) reactions is OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

In this case, the Fe³⁺ ion is gaining an electron, so it is being reduced.

Question 30

Define:

an electrochemical cell

Answer 30

An electrochemical cell is a set of chemical systems which either:

1. undergo a reaction and generate electric current, or
2. undergo a reaction when electric current is run through the system.

Question 31

Identify the principle components of the electrochemical cell shown below.

Answer 31

An electrochemical cell has two half-cells, each of which has an electrolyte and an electrode. The electrodes are connected via electrically conductive material (often a wire).

The half-cells can be separate, as below, or share an electrolyte, but there will always be two separate electrodes.

Question 32

What is the anode of an electrochemical cell?

Answer 32

The anode is the electrode of the half-cell in which oxidation is taking place.

Question 33

What is the cathode of an electrochemical cell?

Answer 33

The cathode is the electrode of the half-cell in which reduction is taking place.

Question 34

Define:

electrolysis

Answer 34

Electrolysis is the act of using a voltage source to drive a non-spontaneous chemical reaction.

Electrochemical cells in which electrolysis is occurring are referred to as electrolytic cells.

Question 35

In an electrochemical cell, at which electrode does oxidation occur? At which electrode does reduction occur?

Answer 35

In any electrochemical cell, oxidation occurs at the anode, while reduction occurs at the cathode.

A simple mnemonic to remember this by: AN OX, RED CAT
ANode = OXidation, REDuction = CAThode.

Question 36

In an electrochemical cell, in what direction do electrons flow?

Answer 36

Electrons always flow from Anode to Cathode.

Remember: Reduction occurs at the Cathode (RED CAT), and Reduction Is Gain of electrons (OIL RIG), so electrons must be flowing to the cathode to facilitate reduction.

Question 37

What type of electrochemical cell is being depicted below?

Answer 37

The cell is an electrolytic cell.

The two main types of electrochemical cells (electrolytic and galvanic) can be easily discerned by the presence or absence of a power source. In this case, the battery is the power source, indicating an electrolytic cell.

Question 38

In the electrolytic cell depicted below, which electrode is the anode?

Answer 38

Electrode A is the anode.

The positive end of the battery attracts electrons towards itself, and those electrons must be leaving Electrode A. Since electrons flow from anode to cathode, Electrode A must be the anode.

Note: Since the anode is attached to the positive end of the battery, in an electrolytic cell the anode is the positive (+) electrode.

Question 39

In the electrolytic cell depicted below, which electrode is the cathode?

Answer 39

Electrode B is the cathode.

The negative end of the battery rejects electrons away from itself, and those electrons must reach Electrode B. Since electrons flow from anode to cathode, Electrode B must be the cathode.

Note: Since the cathode is attached to the negative end of the battery, in an electrolytic cell the cathode is the negative (-) electrode.

Question 40

In the electrolysis of water shown below, which gaseous species will appear at the anode?

Answer 40

O₂(g) will appear at the anode.

The anode is where oxidation occurs, so look for a species which can lose electrons.

Water can be thought of as consisting of O^2- ions and H⁺ ions. The O^2- ions can lose electrons to form molecular oxygen, O₂, and this will occur at the anode.

Question 41

In the electrolysis of NaCl shown below, which chemical species will appear at the anode?

Answer 41

Cl₂(g) will appear at the anode.

NaCl is made up of Na⁺ and Cl^- ions. Cl^- ions will lose electrons at the anode to form molecular Chlorine, Cl₂.

Question 42

In the electrolysis of water shown below, which gaseous species will appear at the cathode?

Answer 42

H₂(g) will appear at the cathode.

The cathode is where reduction occurs, so look for a species which can gain electrons.

Water can be thought of as consisting of O^2- ions and H⁺ ions. The H⁺ ions can gain electrons to form molecular hydrogen, H₂, and this will occur at the cathode.

Question 43

In the electrolysis of NaCl shown below, which chemical species will appear at the anode?

Answer 43

Na(s) will appear at the anode.

NaCl is made up of Na⁺ and Cl^- ions. Na⁺ ions will gain electrons at the anode to form Na metal.

Question 44

What is an electrolyte, in the context of an electrochemical cell?

Answer 44

An electrolyte is any species added to the cell (in which an electrochemical reaction is occurring) which allows current to flow more easily.

Ex: in the electrolysis of water, either acid or salt must be added to the water to improve current flow, since pure water is not a good conductor of electricity.

Question 45

What are the units of current, I?

Answer 45

The units of current are amperes, A.

Current is defined as charge over time Q/t, so 1 A = 1 C/s.

Question 46

If 2.0 A of current flows for three minutes, how many C of charge flow in this time?

Answer 46

360 C

I = 2.0 A = 2.0 C/s
Q = I x t = 2.0 * 3 min * 60s/min
= 360 C

Question 47

What is the equation relating current to number of moles of electrons?

Answer 47

Faraday’s Equation,
I x t = n x F

where:
I = current in A
t = seconds the current flows
n = # of moles of electrons
F = Faraday’s constant, 96,485 C/mol e^-

Note: for the AP Chem exam, approximate F = 10⁵ C/mol e^-

Question 48

In the electrolytic cell shown below, if a current of 1.0 A flows for 10 seconds, how many grams of Na(s) are plated onto the cathode?

Answer 48

2.3x10^-3 g of Na will be deposited.

Use Faraday’s equation to calculate the number of moles of electrons which flow through the circuit:
n = I x T / F
= (1)10/10⁵ = 1x10^-4 mol e^-

Each mole of electrons reduces one mole of Na^+. Calculate the mass of 10^-4 mol of Na(s)
m = 1x10^-4 * 23g/mol
= 2.3x10^-3 g Na

Question 49

What type of electrochemical cell is being depicted below?

Answer 49

The cell is a galvanic cell.

The two main types of electrochemical cells (electrolytic and galvanic) can be easily discerned by the presence or absence of a power source. In this case, the lack of a battery indicates that this is a galvanic cell.

Question 50

In the galvanic cell depicted below, which electrode is the anode?

Answer 50

The Zn electrode is the anode.

In any electrochemical cell, electrons flow from anode to cathode. Electrons are leaving the Zn electrode, so it must be the anode.

Question 51

In the galvanic cell depicted below, which electrode is the cathode?

Answer 51

The Cu electrode is the cathode.

In any electrochemical cell, electrons flow from anode to cathode. Electrons are arriving at the Cu electrode, so it must be the cathode.

Question 52

What is the charge on the cathode of a galvanic cell?

Answer 52

The cathode has a positive (+) charge in a galvanic cell.

The cathode spontaneously attracts electrons, which is what would be expected of a positively-charged electrode. Note that this is the opposite convention from electrolytic cells, where the cathode has a negative charge.

Question 53

What is the charge on the anode of a galvanic cell?

Answer 53

The anode has a negative (-) charge in a galvanic cell.

The anode spontaneously discharges electrons, which is what would be expected of a negatively-charged electrode. Note that this is the opposite convention from electrolytic cells, where the anode has a positive charge.

Question 54

In the galvanic cell depicted below, is the Zn anode going to get lighter or heavier as current flows through the circuit?

Answer 54

The Zn anode will get lighter as current flows.

Oxidation occurs at the anode in any electrochemical cell. In this case, the oxidation reaction will result in Zn(s) going to Zn²⁺(aq). As the Zn ions move into solution, the Zn electrode will get lighter.

Question 55

In the galvanic cell depicted below, is the Cu cathode going to get lighter or heavier as current flows through the circuit?

Answer 55

The cathode will get heavier as current flows.

Reduction occurs at the cathode in any electrochemical cell. In this case, the reduction reaction will result in Cu²⁺(aq) going to Cu(s). As the Cu²⁺ ions move out of solution, they will plate on the Cu electrode, which will get heavier.

Question 56

In the galvanic cell depicted below, which electrode will have metal plated onto it and get heavier?

Answer 56

The Cu cathode will get heavier as Cu ions plate onto its surface.

Question 57

If the overall reaction of an electrochemical cell is:

Co³⁺(aq) + Na(s) ⇒
Co²⁺(aq) + Na⁺(aq)

what is the oxidation half-reaction?

Answer 57

The oxidation half-reaction is:

Na(s) ⇒ Na⁺(aq) + e^-

An oxidation half-reaction follows only the species that gets oxidized in a redox rection, and includes the electron or electrons that flow, as well.

Question 58

If the overall reaction of an electrochemical cell is:

Co³⁺(aq) + Na(s) ⇒
Co²⁺(aq) + Na⁺(aq)

what is the reduction half-reaction?

Answer 58

The reduction half-reaction is:

Co³⁺(aq) + e^- ⇒ Co²⁺(aq)

A reduction half-reaction follows only the species that gets reduced in a redox reaction, and includes the electron or electrons that flow, as well.

Question 59

Define:

the reduction potential of an electrochemical cell Eº_red

Answer 59

Eº_red is the amount of voltage either required or liberated by the reduction half-reaction.

The more positive a species’ reduction potential, the more strongly it will tend to be reduced in electrochemical cells.

Question 60

If the potential of the half reaction:

Co³⁺(aq) + e^- ⇒ Co²⁺(aq)

is Eº_red = +1.82V, what is the potential of a cell which is reducing Co³⁺to Co²⁺?

Answer 60

Eº_red = +1.82 V

In general, the reduction potential gives the amount of voltage released or required by a cell for the reduction reaction to occur. Since the Eº_red > 0 for this cell, it is favorable to reduce Co³⁺, and energy is released when it happens.

Question 61

If the potential of the half reaction:

Co³⁺(aq) + e^- ⇒ Co²⁺(aq)

is Eº_red = +1.82V, what is the potential of a cell which is oxidizing Co²⁺to Co³⁺?

Answer 61

Eº_ox = -1.82 V

An oxidation reaction is the reverse of a reduction reaction; in this case, the reverse of the reduction of Co³⁺. The oxidation potential Eº_ox is simply the reverse of Eº_red. Since Eº_ox < 0, the oxidation is unfavorable to occur, unless 1.82 V are supplied to oxidize this system.

Question 62

If the oxidation potential of a redox half-reaction is Eº_ox, and the reduction potential is Eº_red, what is the** (total) full potential** of the cell?

Answer 62

The cell’s full potential is:

Eº_cell = Eº_ox + Eº_red

Remember to keep track of the appropriate signs of the oxidation and reduction half-reactions; a classic way that chemistry tests will try to trick you is to give you a reduction potential, then ask about the oxidation potential for the same reaction.

Question 63

What is the sign of the potential of a galvanic cell?

Answer 63

Eº_cell > 0 for any galvanic cell.

Since a galvanic cell is spontaneous, it must have a favorable voltage, which means it is positive in sign.

Question 64

What is the sign of the potential of an electrolytic cell?

Answer 64

Eº_cell < 0 for any electrolytic cell.

Since an electrolytic cell is non-spontaneous, it must have an unfavorable voltage, which means it is negative in sign.

Question 65

If the following two reactions are combined into a single galvanic cell, what is the cell’s total potential?

Co³⁺(aq) + e^- ⇒ Co²⁺(aq) Eº = 1.82 V
Na⁺(aq) + e^-⇒ Na(s) Eº = -2.71 V

Answer 65

The cell’s total potential is +4.53 V.

Since the cell is galvanic, its potential must be positive. So the two reactions must be combined (one oxidation, one reduction) to have an overall positive cell potential. In this case, that means that Co must be reduced (Eº_red = +1.82V) while Na is oxidized (Eº_ox = -(-2.71V)).

So the total potential is
Eº_cell = Eº_red(Co) + Eº_ox(Na)
= 1.82 + 2.71 = 4.53 V

Question 66

If the following two reactions are combined into a single electrolytic cell, what is the cell’s total potential?

Co³⁺(aq) + e^- ⇒ Co²⁺(aq) Eº = 1.82 V
Na⁺(aq) + e^- ⇒ Na(s) Eº = -2.71 V

Answer 66

The cell’s total potential is -4.53 V.

Note that the answer is exactly the inverse of these two cells arranged as a galvanic cell. This must be, since the electrolytic cell is just the galvanic cell run backwards. So Co is being oxidized, while Na is being reduced, and the total potential is:

Eº_cell = Eº_red(Na) + Eº_ox(Co)
= -2.71 + (-1.82) = -4.53 V

Question 67

What is the equation relating an electrochemical cell’s standard potential Eº to its standard Gibbs’ Free Energy ΔGº?

Answer 67

ΔGº = -nFEº

where
n = # of moles of electrons in the equation
F = 10⁵ C/mol electrons

Question 68

According to the equation

ΔGº = -nFEº

what is the ΔGº of an electrolytic cell?

Answer 68

For an electrolytic cell, ΔGº > 0

The standard voltage Eº of an electrolytic cell is negative, which means ΔGº must be positive. This can also be derived from the fact that an electrolytic cell is non-spontaneous, and any non-spontaneous reaction has a positive ΔGº.

Question 69

According to the equation

ΔGº = -nFEº

what is the ΔGº of a galvanic cell?

Answer 69

For a galvanic cell, ΔGº < 0

The standard voltage Eº of a galvanic cell is positive, which means ΔGº must be negative. This can also be derived from the fact that a galvanic cell is spontaneous, and any spontaneous reaction has a negative ΔGº.

Question 70

What is the equation relating the standard voltage of an electrochemical cell Eº to its equilibrium constant k_eq?

Answer 70

nFEº = RT ln(k_eq)

Where:
R = ideal gas constant in J/K*mol
T = temperature in K
n = # of moles of electrons
F = 10⁵ C/mol electrons

Question 71

According to the equation

nFEº = RT ln(k_eq)

what is the value of k_eq for a galvanic cell?

Answer 71

For a galvanic cell, k_eq > 1.

For a galvanic cell, Eº is positive, so ln(k_eq) must be positive, hence k_eq > 1. This can also be derived from the fact that a galvanic cell is spontaneous, and for any spontaneous reaction, k_eq > 1.

Question 72

According to the equation

nFEº = RT ln(k_eq)

what is the value of k_eq for an electrolytic cell?

Answer 72

For an electrolytic cell, k_eq < 1.

For an electrolytic cell, Eº is negative, so ln(k_eq) must be negative, hence 0\eq< 1. This can also be derived from the fact that an electrolytic cell is non-spontaneous, and for any non-spontaneous reaction, 0\eq< 1.

Question 73

How can the voltage be calculated for an electrochemical cell which is not at standard conditions?

Answer 73

Use the equation: nFE = -RT ln(Q/k_eq)

Note that this is similar to the equation for calculating the standard voltage. In fact, at standard conditions (where Q = 1), this equation becomes the calculation for standard voltage.

Question 74

According to the equation

nFE = -RT ln(Q/k_eq)

what is the voltage of a cell for which Q < k_eq?

Answer 74

If Q < k_eq, E_cell > 0.

If Q < k_eq, the reaction has not yet reached equilibrium, and will move forwards to reach it. So, the reaction will proceed spontaneously, the cell is galvanic, and E_cell will be positive.

Question 75

According to the equation

nFE = -RT ln(Q/k_eq)

what is the voltage of a cell for which Q > k_eq?

Answer 75

If Q > k_eq, E_cell < 0.

If Q > k_eq, the reaction has exceeded equilibrium, and will proceed backwards to reach it. Since E_cell always refers to the forward reaction, E_cell will be negative.

Question 76

According to the equation

nFE = -RT ln(Q/k_eq)

what is the voltage of a cell for which Q = k_eq?

Answer 76

If Q = k_eq, E_cell = 0.

If Q = k_eq, the reaction is at equilibrium, and will hold at its current state. Since neither the forward nor backward reaction will proceed, E_cell = 0.