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Solubility Flashcards

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1
Q

When do solutions form?

A

When the formation of intermolecular bonds between the solvent and solute is more favorable than the intermolecular bonds within the solvent and solute
Polar solutes dissolve in polar solvents and nonpolar solutes dissolve in nonpolar solvents

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2
Q

Oxidation-Reduction (Redox) Reactions

A

Reactions in which electrons are transferred from one species to another
Electrochemistry: study of redox reactions and the energy changes associated with them
- electrochemical cells: chemical energy of redox reactions is converted into electrical energy that can be used to do work and an electrical circuit is established by migration of charge-carrying ions in solution

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3
Q

Solution

A

Homogeneous mixture of two or more compounds in a single phase (solid, liquid, or gas)
E.g. brass is a solid solution of zinc and copper

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4
Q

Solvent

A

Compound in solution of which there is more of

Solvent is said to dissolve the solute

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5
Q

Solute

A

In a solution, the compound of which there is less

Solute is said to dissolve in the solvent

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6
Q

What are polar molecules held together by? What are nonpolar molecules held together by?

A

Polar molecules: strong intermolecular bonds formed by attraction between partially charged ends
- Polar solute interacts strongly with polar solvent by breaking the solvent-solvent bonds and forming solvent-solute bonds
Nonpolar molecules: Held together by weak intermolecular bonds between instantaneous dipole moments (London dispersion forces)
- Nonpolar solute dose not have enough separation of charge to interact effectively with a polar solvent (cannot spread out within solvent), however can tear apart weak bonds of nonpolar solvent

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7
Q

What happens when an ionic compound dissolves in a polar substance?

A

Cations and anions break apart and are surrounded by oppositely charged ends of the polar solvent (solvation)
Water is good solvent for ionic substances
- Point partially positive hydrogen atoms toward the anions and their partially negative oxygen atoms toward the cations
- When several water molecules attach to one side of an ionic compound, can overcome strong ionic bonds and break apart the compound
- Hydration: molecules then surround the ion

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8
Q

Hydration

A

When water molecules surround an ion
- Said to be in an aqueous phase
- Hydration Number: Number of water molecules that must surround an ion for hydration to occur which varies according to size and charge of ion (commonly 4-6)
When ions form in aqueous solution, solution able to conduct electricity

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9
Q

Electrolyte

A

Compound that forms ions in aqueous solution
Can conduct electricity
Strong electrolytes create solutions that conduct electricity well and contain many ions
Weak electrolytes form fewer ions in solution

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10
Q

Nitrate

A

NO2-

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11
Q

Nitrate

A

NO3-

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12
Q

SO3^2-

A

Sulfite

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13
Q

SO4^2-

A

Sulfate

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14
Q

Hypochlorite

A

ClO-

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15
Q

Chlorite

A

ClO2^-

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16
Q

Chlorate

A

ClO3^-

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17
Q

Perchlorate

A

ClO4^-

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18
Q

Carbonate

A

CO3^2-

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19
Q

Bicarbonate

A

HCO3^-

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20
Q

Phosphate

A

PO4^3-

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21
Q

Ammonium

A

NH4^+

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22
Q

Molarity

A

moles of solute divided by the volume of the solution
Units: usually mol/L
M = moles of solute / volume of solution

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23
Q

Molality

A

Moles of solute divided by kilograms of solvent
Units: mol/kg
Used to calculate freezing point depression and boiling point elevation
m = moles of solute / kilograms of solvent

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24
Q

Mole fraction

A

Moles of a compound divided by the total moles of all species in solution
Ratio, so mole fraction has no units
Used to calculate vapor pressure of a solution
\chi = moles of solute / total moles of all solutes and solvent

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25
Q

Mass percentage

A

Ratio of the mass of the solute total mass of solution multiplied by 100
Mass % = ( mass of solute / total mass of solution) x 100%

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26
Q

Parts per million

A

10^6 multipled by the ratio of the mass of solute to the total mass of the solution
Ppm = (mass of solute / total mass of solution) x 10^6
NOT the number of solute molecules per million molecules

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27
Q

What are the five ways to represent solute concentration?

A

Molarity, molality, mole fraction, mass percentage, parts per million, Normality
Always given in terms of the form of solute before dissolution
E.g. 1 mol of NaCl added to 1 L of water, resulting solution is ~ 1 molar, not 2 molar even though one mole of NaCl dissociates into two moles of ions

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28
Q

Normality

A

Number of equivalents per liter of solution
Equivalent: depends on type of reaction taking place in solution
- Acid-base reaction: equivalent is defined as mass of acid or base that can donate or accept one mole of protons
- 1 molar solution of H2SO4 is called a 2 normal solution because it can donate 2 protons for each H2SO4 molecule

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29
Q

Steps of forming a solution

A

Physical reaction involving three steps:

  1. Breaking of intermolecular bonds between solute molecules
  2. Breaking of intermolecular bonds between solvent molecules
  3. Formation of intermolecular bonds between solvent and solute molecules
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30
Q

What is the enthalpy change of forming a solution (not at high pressure)?

A

Enthalpy change is approximately equal to internal energy change
\delta H ~= \delta U
Heat of solution (\delta H_sol) = \delta H1 + \delta H2 + \delta H3
First two steps are endothermic and third step is exothermic
Hesse’s Law: overall enthalpy change of a reaction can be exothermic even if the reaction contains endothermic steps

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31
Q

Is the solution with new intermolecular bonds more stable if the reaction of forming the solution is endothermic or exothermic?

A

More stable if the reaction was exothermic (released heat)
new intermolecular bonds are more stable than original ones and new intermolecular attractions within the solution stronger than intermolecular attractions within pure substances

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32
Q

What does the heat of hydration refer to in the formation of a solution?

A

Combining steps 2 and 3 of the formation of a solution, which is the breaking of intermolecular bonds between solvent molecules and the formation of intermolecular bonds between solvent and solute molecules

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33
Q

What is the change in entropy usually for the formation of a solution?

A

Usually involves an increase in entropy of system

One exception is when a gas dissolves in a liquid or solid (entropy change will be negative)

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34
Q

How does adding heat sometimes help dissolve a solute in a solvent?

A

Adding heat gives extra energy to break intermolecular bonds between solute-solute and solvent-solvent to form new bonds
Without input of heat, \delta G would remain positive and dissolution would not happen at any appreciable level

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35
Q

What can the Gibbs energy equation tell us about when the formation of a solution will proceed spontaneously?

A

Gibbs free energy equation is: \delta G = \delta Hsol - T \delta S
Entropy is usually positive for dissolution, but \delta Hsol can be positive or negative
If \delta Hsol is negative, then will proceed spontaneously
If \delta Hsol is positive, then temperature dependent

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36
Q

If there is a pure liquid in a vacuum-sealed container, is the space above the liquid a vacuum?

A

Certain amounts of liquid molecules escape the liquid intermolecular bonds with enough kinetic energy and go into the gas phase
So there are some gas molecules above the liquid
As the space above the liquid fills with molecules, some of the molecules crash back into the liquid
Equilibrium reached when rate at which molecules are leaving the liquid equal rate at which molecules are re-entering liquid

37
Q

Vapor Pressure

A

Partial Pressure of the compound that is required to create equilibrium of liquid molecules going into gas and gas molecules going into liquid
- created by molecules in the open space at equilibrium
Related to kinetic energy of the molecules, so a function of temperature

38
Q

What happens to the partial pressure of water vapor above a puddle and the pressure of the water in the puddle when wind blows over the puddle?

A

Partial pressure of water vapor decreases when the wind blows (velocity increases, so pressure decreases)
Pressure of water does not change, since it is fixed at a given temperature
Since partial pressure of water vapor now lower than pressure of water (they were in equilibrium), number of water molecules leaving the puddle is greater than number of molecules entering the puddle
Water is evaporating from puddle

39
Q

When a liquid is boiling, what does this mean about it’s vapor pressure wrt the local atmospheric pressure?

A

When vapor pressure of liquid equals local atmospheric pressure, liquid boils
A liquid can be brought to a boil by raising the temperature, thereby raising the vapor pressure until it reaches the atmospheric pressure OR by lowering the atmospheric pressure until it equals the vapor pressure

40
Q

Do solids have vapor pressure?

A

Solids do have a vapor pressure, usually very low
Melting point is temperature at which the vapor pressure of solid is equal to vapor pressure of liquid phase of that substance
Above the melting point, the vapor pressure of the liquid is _ than that of the solid, so solid moves into liquid

41
Q

Nonvolatile Solute

A

Solute with no vapor pressure
If added to a liquid, some solute molecules reach the surface of the solution and reduce the surface area available for liquid molecules
Since solute molecules do not break free form solution, but take up surface area, number of molecules breaking free from liquid decreases while surface area of solution and volume of open space above the solution remain the same
- Means that vapor pressure of solution is lower than vapor pressure of pure solvent

42
Q

Raoult’s Law

A

Pv = \chi_a Pa

Vapor pressure of solution Pv is proportional to vapor pressure of pure liquid Pa and the mole fraction of liquid \chi_a

43
Q

Volatile solute

A

A solute with a vapor pressure
Will also compete for the surface area of a liquid, however some of the molecules of a volatile solute will escape form the solution and contribute to vapor pressure
For ideal solution, in which the solute and solvent have similar properties, partial pressures contributed by solvent and solute can be found by applying Raoult’s law separately: Pv = \chi_a Pa + \chi_b Pb

44
Q

How will the vapor pressure of a nonideal solution deviate from that predicted from Raoul’s law?

A

Can predict direction of deviation based on heats of solution
If heat of solution is negative, stronger bonds are formed and fewer molecules are able to break free from surface, so negative deviation of vapor pressure from Raoult’s law
Opposite for positive heat of solution (endothermic)

45
Q

Solubility

A

Quantifies a solute’s tendency to dissolve in a solvent
Maximum number of moles of a solute that can dissolve in a solution (depends on temperature and the ions in solution)
Commonly measured in mol/L
If a salt, dissolution is reversible and at first the rate of reverse reaction, precipitation, is lower than rate of dissolution
As conc. Of dissolved salt increases, rates of dissolution and precipitation equilibrate (solution is said to be saturated)

46
Q

Solubility Product

A

Ksp where solids and pure liquids are left out of the solubility product expression
Computed the same way as other equilibrium constant, K
Concentration of products over concentration of reactants, each raised to the power of their coefficients in the balanced equation
Fixed for a given temperature, independent of ion concentrations
Measure of how inclined a compound is to dissolve

47
Q

How can you find solubility from the solubility product?

A

Solubility: # moles of solute / L of solution that can be dissolved in a given solvent
For dissolution of BaF2(s) into Ba^2+(aq) + 2F^-(aq), Ksp = [Ba^2+][F^-]^2
At 25 deg C, Ksp = 2.4 x 10^-5
Set x = # moles / L OR solubility
2.4 x 10^-5 = (x)(2x)^2 from Ksp equation
Where x ~= 1.8 x 10^-2 mol / L for 1 L of water at 25 deg C

48
Q

Common Ion Effect

A

Disturbance of the equilibrium of solubility of an ion that is the same as one that is already in solution
E.g. adding NaF to a solution that already contains BaF2 dissolved in equilibrium
Pushes equilibrium in direction that will reduce the concentration of that ion
Spectator Ion: Na+ would be a spectator, since it does not affect the equilibrium

49
Q

In the case where we dissolve NaF in 1L of water to completely dissociate and form 1 mol of F- and 1 mole of Na+ and then we add BaF2 to this solution, how do we calculate the solubility of F- in BaF2?

A

If Ksp = 2.4 x 10^-5 for 25 deg C adn we are at 25 deg C:
2.4 x 10^-5 = (x)(2x + 1)^2
Can simplify math, because since x will be very small, can say that 2x + 1 ~= 1
2.4 x 10^-5 = (x)(1) -> x = 2.4 x 10^-5
Double check if 2x is much less than 1
4.8 x 10^-5 &laquo_space;1
A common ion added to a solution that is not saturated does NOT shift the equilibrium, because in unsaturated solution there is no equilibrium to shift

50
Q

Solubility of ionic compounds containing nitrate (NO3^-), ammonium (NH4^+), and alkali metals (Li+, Na+, K+,…)?

A

Soluble!

51
Q

Solubility of ionic compounds containing halogens?

A

(Cl-, Br-, I-) are soluble

Mercury, lead, and silver compounds are not soluble (Hg2^2+, Pb^2+, Ag^+)

52
Q

Solubility of sulfate compounds?

A

Sulfate: SO4^2-, most are soluble
Except for those containing mercury, lead, and heavier alkaline earth metals (Hg2^2+, Pb^2+, Ca^2+, Sr^2+, Ba^2+) -> not soluble

53
Q

Solubility of compounds containing the heavier alkaline metals?

A

Ca^2+, Sr^2+, Ba^2+ are soluble when paired with sulfides (S^2-) and hydroxides (OH^-)

54
Q

Solubility of carbonates, hydroxides, phosphates, and sulfides?

A

CO3^2-, PO4^3-, S^2-, OH^- are generally insoluble other than OH- and S^2- with heavier alkaline metals

55
Q

What is the cutoff for consideration of soluble vs. insoluble generally?

A

If solubility less than 0.01 mol / L, than generally said to be insoluble
Ions with single charges are generally soluble, and ions with double charges on one ion are less soluble
Smaller ions or molecules such as those in Group 1 are soluble
Large, heavy cations are less soluble
Higher charges means the molecule might have stronger intramolecular ionic bonds, so individual ions are harder to pull apart and dissolve

56
Q

Henry’s Law

A

C = ka1 Pv
C is solubility of gas a (mol / L), ka1 is Henry’s law constant (unique to each solute-solvent pair), Pv is vapor partial pressure of gas a
Concentration of gas in solution is proportional to vapor partial pressure of gas above solution
Pressure and temperature can affect solubility (pressure increases solubility of gas)
Law can also be written as:
Pv = \chi_a ka2, where \chi_a is mole fraction of a in solution, and ka2 is Henry’s law constant

57
Q

What is the conflict between Henry’s Law and Raoult’s Law?

A

In the second formulation of Henry’s Law: Pv = \chi_a ka2
And Raoult’s Law: Pv = \chi_a Pa
But, ka2 does NOT equal Pa, so they are in conflict
Both are approximations and are more accurate in certain contexts
Raoult’s law is more accurate when it is applied to vapor partial pressure of a solvent with high concentration
Henry’s law is more accurate when applied to vapor partial pressure of a volatile solute where the solute has a low concentration
In an ideally dilute solution, the solvent obeys Raoult’s law and the solute obeys Henry’s law
When solvent conc is high, each solvent molecule is surrounded by other solvent molecules, so behaves more like pure solvent, more similar to vapor pressure of pure liquid (Raoult’s)

58
Q

What is a caution to have when using Le Chatelier’s principle to predict what temperature will push the reaction left or right?

A

In an exothermic dissolution reaction, heat will be produced
Le Chatelier’s principle tells us that more heat should push the reaction to the left, however it does not
The dissolution is accompanied by a large increase in entropy which in \delta G = \delta H - T \delta S means that increased temperature can increase spontaneity of dissolution
Due to large increase in entropy, water solubility of many solids increases with increasing temperature regardless of the enthalpy change (must be found by experiment to be absolutely certain)
Not the same when a gas dissolves in a liquid, bc entropy change is negative for this dissolution. So as temperature increases, gas solubility generally decreases

59
Q

Why is hot waste water from factories that is dumped into streams hazardous to aquatic life?

A

Double effect

  1. Hot water holds less oxygen than cold water (gas solubility decreases with increasing temperature)
  2. Hot water floats on the cold water, sealing it off from oxygen in the air above
60
Q

How do acids and bases affect solubility?

A

For example, if we had: Ba(OH)2 (s) -> Ba^2+(aq) + 2OH^- (aq)
If we add base to the solution, it will force reaction to the left (Le Chatelier)
If we add acid, then we have H^+ (aq) + OH-(aq) -> H2O(l) which removes the hydroxide ions from the solution, and more of the solid compound would be dissolved to maintain equilibrium

61
Q

What are some factors that effect the solubility of a gas?

A

Increased temperature of solvent decreases solubility
Heavier, larger gases experience greater Van der Waals forces and tend to be more soluble
Gases that chemically react with solvent have greater solubility

62
Q

Redox reactions

A

Electrons are transferred from one atom to another

Atom that loses electrons is oxidized, atom that gains electrons is reduced

63
Q

Oxidation states

A

Possible charge values that an atom can hold within a molecule
In many cases, the charges do not truly exist, and oxidation states simply provide a system for tracking the movement of electrons
Oxidation states must add up to total charge on molecule or ion
Often the same as the charge of the cation or anion that the element typically forms (difference to Noble gas config), but if atoms had permanent oxidation states, redox reactions could not take place

64
Q

General Oxidation State Rules

A

Rules higher on table have higher priority if contradictions:
Oxidation State Atom
0 Atoms in their elemental form
-1 Fluorine
+1 Hydrogen (except when bonded to a metal, like NaH; then -1)
-2 Oxygen (except when in a peroxide like H2O2; then -1)

65
Q

Group Oxidation States

A

Rows higher up in table have higher priority during contradictions
Oxidation State Group on Periodic Table
+1 Group 1 elements (alkali metals)
+2. Group 2 elements (alkaline earth metals)
+3 Group 15 elements (Nitrogen family)
-2 Group 16 elements (Oxygen family)
-1 Group 17 elements (halogens)

66
Q

Reducing agent and oxidizing agent

A

Reducing agent (AKA reductant) gives electrons to another species, becomes oxidized
Oxidizing agent (AKA oxidant) receives electrons from another species, so becomes reduced and oxidizes the other species
E.g. CH4 + 2O2 -> CO2 + 2H2O
C has gone from -4 to +4, Oxygen goes from 0 to -2
Atom is oxidized or reduced, compound is the oxidizing or reducing agent

67
Q

Redox Titration

A

Used to find the molarity of a reducing agent (remember, mol / L)
Strong oxidizing agent is titrated and the resulting voltage change is measured
Solutions must be different, one is standard solution and one has reducing agent
As strong oxidizing agent is added to solution with reducing agent, voltage increases first gradually and then quite suddenly
Half equivalence point: point near the middle of the gradual increase
Equivalence point: voltage suddenly shoots up
- Occurs when all moles of reducing agent in solution have been completely oxidized
- # moles of oxidizing agent required to reach equivalence point will be either equal to or a multiple of the number of moles of the reducing agent in the solution that is being titrated (one molecule of oxidizing agent may accept a different number of electrons than one molecule of reducing agent gives up)

68
Q

How can concentration of Sn^2+ in 100 mL solution be determined using a redox titration?

A

For example, if you add oxidizing agent Ce^4+, will take electrons from Sn^2+ in following two reactions:
Sn^2+ -> Sn^4+ + 2e-
Ce^4+ + e- -> Ce^3+
Where two Ce^4+ atoms are required for each Sn^2+ atom
If add 2mL of 5 mM Ce^4+ solution to get to equivalence point, molarity of original Sn^2+ solution is determined by:
2 mL Ce^4+ solution x 5 mmol / L x 1L / 1000 mL = 0.01 mmoles of Ce^4+
This means 0.01 mmoles Ce^4+ are required to oxidize 0.005 mmoles Sn^2+
Therefore:
(0.005 mmoles Sn^2+ / 100 mL unknown Sn^2+ solution) x (1000 mL / L) = 0.05 mM Sn^2+ solution

69
Q

Electric Potential (E) in Redox reaction

A

Likelihood that the redox reaction will proceed
The more positive the potential, the more likely the reaction will proceed
Potentials for oxidation and reduction components of reaction can be approximated separately based on a standard hydrogen electrode (SHE)
- Each component is a half reaction (can’t occur by itself)
- Reverse of a reduction half reaction is an oxidation half reaction
Half reaction potentials are usually listed as reduction potentials, but reverse sign for oxidation potential

70
Q

Half reaction at a standard hydrogen electrode

A

2H+ + 2e- -> H2
Where E^{o} noted as 0.00 V
All electrical potentials for redox reactions are relative to the Standard hydrogen electrode, so positive means greater than SHE and negative means less than SHE

71
Q

Calculate the potential of 2Au^3+ + 3Cu -> 3Cu^2+ + 2Au if we know:
Au^3+(aq) + 3e- -> Au(s), E^{o} = 1.50 V
Cu^2+(aq) + 2e- -> Cu(s), E^{o} = 0.34 V

A

Reduction potentials are intensive properties, so do not multiply by number of ions in each half reaction
2(Au^3+ + 3e- -> Au), E^{o} = 1.50 V
3(Cu -> Cu^2+ + 2e-), E^{o} = -0.34V
1.50V - 0.34V = 1.16V

72
Q

What are the steps to balance a redox reaction?

A

Balance redox reaction in acidic solution:
1. Divide the reaction into its corresponding half reactions
2. Balance the elements other than H and O
3. Add H2O to one side until the O atoms are balanced
4. Add H+ to one side until the H atoms are balanced
5. Add e- to one side until the charge is balanced
6. Multiply each half reaction by an integer so that an equal number of electrons are transferred into each reaction
7. Add the two half reactions and simplify
For redox reactions occurring in basic solution, follow same steps, then neutralize H+ ions by adding the same number of OH- ions to both sides of reaction

73
Q

Galvanic Cell

A

Cell in which two distinct electrically conducting chemical phases are placed in contact and a charged species from one phase cannot flow freely to another, creating an electric potential
Has an alternative pathway for just the flow of electrons between phases, generating a current from one phase to another in a conversion of chemical energy to electrical energy
Made of a multiphase series of components, no components present in more than one phase, and all phases must conduct electricity, but at least one impermeable to e-
Ionic conductor: phase impermeable to e- carries the current in the form of ions (electrolyte sol’n in form of salt bridge)

74
Q

What are the components of a simple galvanic cell?

A

T-E-I-E’-T
T: terminals (conductors such as metal wires)
E: electrodes (also conductors), anode and cathode
I: ionic conductor (often salt bridge)
When cell is formed, emf is electric potential difference between T and T’

75
Q

Anode and Cathode of Galvanic Cell

A

Anode: marked with a negative sign, where oxidation takes place
Cathode: marked with positive sign, where reduction takes place
Sometimes electrode also refers to electrolyte solution and strip of metal, or can be referred to wholly as a half cell

76
Q

Cell Potential

A

symbol E, also called the electromotive force (emf)
Potential difference between the terminals when they are not connected
Connecting the terminals reduces the potential difference due to internal resistance within the galvanic cell
Drop in emf increases as current increases
Current from one terminal to other defined as moving in opposite direction to electron flow (e- flow from anode to cathode)

77
Q

Standard state cell potential

A

Sum of the standard state potentials of the corresponding half reactions
Cell potential for galvanic cell is always positive (a galvanic cell always has chemical energy that can be converted to work)
Cell potential depends on:
- half reactions
- concentrations of reactants and products
- temperature
Concentrations will always be 1M, so you can use values from table to calculate cell potential

78
Q

What is the cell potential of any electrode used in conjunction with the SHE?

A

Exactly equal to the reduction potential of the half reaction occurring at the other electrode
Many half reaction reduction potentials can be measured using the SHE
SHE has no salt bridge, because both electrodes are in contact with the same solution, for example, producing HCl from H2 and AgCl

79
Q

Salt Bridge

A

Type of liquid junction that minimizes potential difference
When cell contains two difference solutions, liquid junction is required to separate the solutions, but because ions can move across a liquid junction, liquid junctions create additional small potential difference that affects potential of galvanic cell
Salt bridges can be made from an aqueous solution of KCl
Allow movement of ions between solutions without creating a strong extra potential, because K+ ions move toward the cathode at about the same rate that the Cl- ions move towards the anode
Job is to maintain the neutrality of the half cell

80
Q

What happens with leakage at a salt bridge over time?

A

Causes the battery to lose its chemical potential over time, because charge buildup will occur

81
Q

Cell Diagram

A

IUPAC convention to represent a galvanic cell where each phase is listed from left to right, beginning with the terminal attached to the anode and ending with the terminal attached to the cathode
Terminals often left out because they are always the same material ( do not take place in reaction)
Vertical line is placed between phases, double vertical line indicates a salt bridge, dotted vertical line indicates a boundary between two miscible liquids, and species in the same phase are separated by a comma
Pt’(s)|Zn(s)|Zn^2+(aq)||Cu^2+(aq)|Cu(s)|Pt(s)

82
Q

How do you calculate the Gibbs free energy for a galvanic cell?

A

\delta G = -nFE_max
Where n is number of moles of electrons transferred in balanced redox reaction, F is charge on one mole of e- (96,486 C / mol), and E_max is cell potential
Says that free energy is the product of total charge nF and voltage E (product of charge (C) and voltage (V) is equal to electrical work)
Positive cell potential means spontaneous reaction
When all conditions are standard:
\delta G^{o} = -nFE^{o}

83
Q

How can potential be found when the concentrations are not one molar?

A

\delta G = \delta G^{o} + RT ln(Q) and substitute -nFE for \delta G and -nFE^{o} for \delta G^{o}, and then divide by -nF:
Nernst Equation: E = E^{o} - (RT / nF) ln(Q)
At 298K and base 10 log:
E = E^{o} - 0.06/n log (Q)
So calculate Q first by plugging in nonstandard concentration to compute the equilibrium constant

84
Q

Concentration Cell

A

Limited form of galvanic cell in which a reduction half reaction takes place in one half cell, while the exact reverse of that half reaction is taking place in the other half cell
Cells differ in ion concentrations
Potential: E^{o} = 0V
If concentrations were equal, concentration cell potential would be zero
Nernst equation can be used to calculate potential for concentration cell, but can determine flow of cell if think about entropy wanting to equalize the concentrations on both sides of cell

85
Q

What is the Nernst equation adapted for a concentration cell?

A

Remember, that the solution in the reaction is both a reactant and a product, so will be in numerator and denominator for different half cell concentrations:
E = E^{o} - 0.06 / 2 log(products/ reactants) with E^{o} = 0V

86
Q

Electrolytic Cell

A

A type of galvanic cell in which there is a power source connected to wire connecting electrodes which can drive the reaction in the opposite direction
Will always have a negative emf
Cathode is marked negative and anode is marked positive
However the reduction still takes place at the cathode and oxidation takes place at the anode
Used in industry for metal plating, purifying metals

87
Q

Comparing reduction potentials of half reactions, how can you tell that a reaction will not run spontaneously?

A

E.g. Na+ + e- -> Na, E^{o} = -2.71V
2Cl- -> 2e- + Cl2, E^{o} = -1.36V
Reaction will not run in aqueous solution, because water has a less negative reduction potential than sodium, which means sodium will oxidize spontaneously in water

88
Q

Lead Storage battery

A

Cell of choice for cars because of high current it can supply
Can act as both a galvanic and electrolytic due to rechargeable capabilities
Anode made of lead, cathode made of lead dioxide
Discharges in the presence of sulfuric acid
Oxidation of lead at anode: Pb(s) -> 2e- + Pb^2+(aq)
Sulfuric acid reacts with lead ions to precipitate insoluble lead sulfate:
Pb^2+(aq) + SO4^2-(aq) -> PbSO4(s)
Reduction at the cathode: 2e- + PbO2(s) + 4H+(aq) + SO4^2-(aq) -> PbSO4(s) + 2H2O(l)
Both half reactions produce lead sulfate, insoluble toxin

89
Q

Nickel Cadmium battery

A

More expensive than lead storage battery
Improvement upon lead storage battery due to size (smaller) and because its electrolyte is not consumed during discharge
Batteries arranged in cylindrical configurations that reduce internal resistance

MCAT Chemistry (8 decks)

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