11th2nd: APCHEM Equilibrium

What do you not include in the Haber equilibrium process?

solids and liquids

Show the relationship between Kc and Kp using PV \= nRT part 1

Isolate P: P \= (nRT)/(V)

Show the relationship between Kc and Kp using PV \= nRT part 2

Determine what represents concentration in the new equation; n/V \= concentration

Show the relationship between Kc and Kp using PV \= nRT part 3

Sub in C (concentration) for n/V: P \= C * RT

Why can we express pressure in terms of concentration and the other way around?

As long as temperature is held constant, pressure and concentration are proportional (P = C \* RT)

For equation (N2 (g) + 3H2 (g)

Kc \= (NH3)^2 ÷ ((H2)^3 * (N2))

For equation (N2 (g) + 3H2 (g)

If Pressure \= Concentration * RT, to get ((P NH3)^2 ÷ ((P H2)^3 * (P N2)), substitute each P term with their corresponding Kc term multiplied by RT

For equation (N2 (g) + 3H2 (g)

((NH3 * RT)^2 ÷ ((H2 * RT)^3 * (N2 * RT))

For equation (N2 (g) + 3H2 (g)

TOP: (NH3)^2 * (RT)^2
BOTTOM: (H2)^3 * (RT)^3 * (N2) (RT)

For equation (N2 (g) + 3H2 (g)

((NH3)^2 / (H2)^3 * (N2)) times (RT)^2 / (RT)^4

For equation (N2 (g) + 3H2 (g)

Kc * (RT)^-2

Kp \=

Kc(RT)^∆n

what does ∆n and T mean in the Kp equation?

∆n = non solid/liquid moles product - non solid/liquid moles reactant

T = absolute temperature

What is the relationship of the K values between the forward and reverse reactions?

The K's are inverses of each other

What happens to K if all gaseous and/or aqueous substance concentrations are doubled?

Since coefficients represent exponents in the expression, doubling concentrations would square K

What happens to K if all gaseous and/or aqueous substance concentrations are halved?

Since coefficients represent exponents in the expression, halving concentrations would raise K to the 1/2

For the reaction CaCO3(s)

Since Kp only changes with temperature and Kp solely \= CO2, then a change in volume or concentration of CaCO3 will not affect Kp, and thus PCO2

What do you have to note about the proportions inside ICE tables?

stoichiometric proportions only apply to the "change" row

What is the relationship between reaction quotient (Q) and equilibrium constant (K)?

If Q is greater than K, then that means there are too many products; if Q is less than K, then there are too many reactants

Why does changing the volume of a container cause a shift in equilibrium position?

Since K takes into consideration concentration (moles/liters), changing the volume would offset the mole to liter proportion, meaning the reaction must shift in order to reach the same proportions again to equal K

Whenever a system at equilibrium experiences a disturbance, as changes occur to return to equilibrium, what do you have to note about the lines of each component?

all graph lines show a decelerating curve if plotted as time on the x axis, because reactions naturally tend to slow down over time

n vs iso prefix for organic compound nomenclature

n = all carbons are in an unbranched, linear chain

iso = all carbons but one are in an unbranched, linear chain

Arrhenius acids and bases

ACIDS increase \[H+\] / \[H3O+\] in water

BASES increase \[OH-\] in water

Arrhenius acids and bases only include …… substances while Brønsted Lowry acids and bases can include …..

\-aqueous

\-aqueous, liquid, solid, gas

How can Brønsted-Lowry acids/bases also be Arrhenius acids/bases?

\
EX: HCl + H2O → H3O+ + Cl-

In water, HCl’s dissociated proton bonds with water, leaving Cl- as the conjugate base and H3O+ as the conjugate acid

6 Strong Acids

HCl, HBr, HI, H2SO4, HClO4, HNO3

What categories of bases are strong?

\-alkali with OH

\-alkalines with (OH)2

Why is the dissociation of a weak acid into water reversible? (applies to weak bases too)

H3O+ and a stronger conjugate base is formed through a weak acid, and OH- and a stronger conjugate acid is formed through a weak base, both species being strong enough to react back

With the dissociation of weak acids/bases, the equilibrium position favors…

the side with the weaker acid/base because those substances have lower potential energy

What are the two main categories of weak bases?

1) neutral atoms that have an atom with a lone pair that can accept a H+

2) anions derived from weak acids

Amines are…

NH3 derivatives with a lone pair on N and at least 1 H replaced with an organic group represented as R

Anions derived from STRONG acids will have (stronger/weaker) basicity than water, so the anion (will/will not) react further and will result in a (neutral/basic/acidic) solution

weaker; will not; neutral

Anions derived from WEAK acids will have (stronger/weaker) basicity than water, so the anion (will/will not) react further and will result in a (neutral/basic/acidic) solution

stronger; will; basic

Cations derived from STRONG bases will have (stronger/weaker) acidity than water, so the anion (will/will not) react further and will result in a (neutral/basic/acidic) solution

weaker; will not; neutral

Cations derived from WEAK bases will have (stronger/weaker) acidity than water, so the anion (will/will not) react further and will result in a (neutral/basic/acidic) solution

stronger; will; acidic

If the cation and anion are derived from a weak base and weak acid, respectively, how is the resulting acidity level of the solution determined?

depends on the Kb and Ka values of the cation and ion

\-ACID = Ka > Kb

What are 3 key factors that make an acid strong?

1) the bond between H and the other ion is highly electronegative and thus polarized

2) weaker bond

3) the resultant anion is more stable

Why does high polarization factor into a strong acid?

the species bonded to H+ has greater ability to draw electrons away from H+, facilitating H+ breaking away and dissociating

Down a group, acid strength tends to (increase/decrease) because …

increase, the electronegativities stay relatively the same while the bond strength will decrease due to additional shielding

Why is HF considered a weak acid in comparison to other H-Halogen acids?

\-since F- is smaller than many other ions in its group, its additional electron unstabilizes it much more severely

In oxoacids, hydrogen is bonded to…

oxygen

How can the central atom and additional atoms/groups of atoms increase an oxoacid’s strength?

Greater electronegativity of the central atom and more, additional atoms increase polarization, further pulling away electrons from H+

Why is HOI acid weaker than HOCl acid?

Since hydrogen in both is bonded to the same species, oxygen, the additional atoms around oxygen determine the acid strength, so iodine’s smaller electronegativity creates less polarity

When asked to solve for the pH, what do you assume about the environment and state of reaction?

unless specified otherwise, always conclude that the problem is asking for pH at equilibrium at 25˚ or 298K

Relationship between pH and pOH?

pH + pOH = 14

As the concentration of H+/OH- ions increases, pH/pOH ... because

decreases, because in a negative log function, y values infinitely get smaller with increasing x values

Conceptually, why does it make sense for more H+ to decrease pH and more OH- to decrease pOH?

more H+ in a solution means higher acidity, so the pH goes down, while more OH- decreases the pOH, so when subtracted from 14, the pH is high to reflect a base

At 25˚C, what are the concentrations of OH- and H3O+?

\[1.0 x 10^-7\]

For reactions involving the dissociation of strong acids or bases, how do you determine the starting concentration of H3O+ or OH-?

H3O+/OH- starting concentration is technically 10^-7, but since if the starting concentration of the acid'/base is significantly larger than ^-7, then approximate the H3O+/OH- concentration as 0 because 10^-7 would be negligible at that point

For reactions involving the dissociation of weak acids or bases, how do you determine the starting concentration of H3O+ or OH-?

assume 0

When can you assume that the minus x change of a weak acid/base reaction is negligible?

since K = products over reactants, a very small K value indicates little dissociation, meaning the minus x is most likely negligible

Percent dissociation equation

(concentration of acid/base dissociated ÷ original concentration) \* 100

What does “p” indicate? Examples?

p tells us to take the -log of a quantity; H, OH, Ka, Kb

A strong acid has a (small/large) Ka value and a (small/large) pKa value

large; small

Kw value at 25C˚

1\.0 x 10^-14

Ka and Kb relationship

Ka x Kb = Kw at 25C

Explain the common ion effect

In an equilibrium system with a weak conjugate acid-base pair, if a strong electrolyte containing one of the conjugates is added, the equilibrium shifts to relieve the disruption, also increasing/decreasing the H+ or OH- concentration from its original amount

In the system “HC2H3O2 + H2O <> H+ + CH3COO-”, what happens if you add NaCH3COO?

Na+ is just a spectator, while the extra CH3COO- will cause a leftward shift, decreasing the H+ and increasing the pH

What is formed between a strong acid and weak base?

very weak conjugate base and a conjugate acid of stronger acidity than H2O

What is formed between a weak acid and strong base?

H2O, a very weak conjugate acid, and a conjugate base of stronger basicity than H2O

What happens between a weak acid and weak base?

Equilibrium occurs; while the weak acid and base are dissociating, their stronger conjugates have enough strength to react back to their original acid base

What are 3 characteristics of buffers?

\-solutions of a weak conjugate acid-base pair

\-both of \~high concentrations

\-both of \~equal concentration

In the BCA table, use …. but in the ICE table, use …..

moles; concentration

When is it ok to use moles in the ICE table (2)?

1) when the total volume stays constant

2) volume happens to be 1 L

How can strong acids and bases be used to form buffers? (2)

1) Titrate a weak base with a strong acid, so that once the reaction completes, \~equal and high amounts of weak acid-base remain

2) Repeat the same steps as #1, but swap the roles of acid and base

After titrating the buffer with a strong acid or base, what do you have the choice of doing next?

When plugging the amounts into the equilibrium expression, you can either choose the dissociation of the weak acid or dissociation of the weak base expression

In a buffer system, what equilibrium reactions are happening simultaneously?

the dissociation of the acid with water into the conjugate base and H3O+, and the dissociation of the base with water into the conjugate acid and OH-

What happens if you add a weak acid or base with a common ion to one of the species in the buffer system?

Unlike adding a strong acid/base, you cannot assume that the added weak base/acid completely dissociates, so just convert it to concentration and put it in the equilibrium expression

Calculating pH after equivalence point shortcut; EX:

HCl + NaHCO3 → Na+ + Cl- + H2CO3

when HCl dissociates entirely into its conjugate and H+, while some acid H2CO3 is still formed and can be plugged into the equilibrium expression, the H+ formed from the equilibrium is insignificant to the H+ formed from the excess HCl

Before approaching the equivalence point, why is the titration curve much flatter for __**strong acid-base compared**__ to the gradual increase of weak acid-strong base?

Any time a strong acid and base react, the result is a neutralized solution with a salt and water, allowing the pH to stay \~constant

Before approaching the equivalence point, why is the titration curve for __**weak acid-strong base**__ more of a gradual increase compared to the flatness of strong acid-strong base?

Since HF is a weak base, when it reacts with NaOH, HF produces F- which can react further with H2O to produce OH-, thus gradually increasing the pH

The length of the x axis before reaching the equivalence point is dependent on what? Why?

the analyte’s unknown concentration as well as the molarity of the titrating solution, because equivalence point is not dependent on acid or base strength but depends on stoichiometric moles reacting

Significance of half equivalence point of a titration? (applies to strong-weak titrations)

Even though the whole initial stage of the titration acts as a buffer, the half equivalence point best acts as one because the added acid/base will cause the analyte base/acid and its conjugate to be of \~equal concentrations

Henderson Hasselburg Equation

pH = pKa + log( \[A-\] / \[HA\] )

What is the x intercept of the mother log function?

(1, 0)

When is pKa = pH

at half equilibrium, because the concentrations of conjugate acid and base will be roughly the same, meaning the part of the Henderson-Hasselburg equation = 0

How do indicators work?

color depends on the H+ concentration, meaning changing the conjugate acid/base concentrations will shift the reaction away or towards H+

If pH is greater than pKa, what does that mean about the concentrations of acid and base?

the base has a greater concentration

What happens when a solution becomes saturated after dissolving enough solid?

Extra solid sinks to the bottom, and equilibrium occurs between the solid and its ions

Ksp vs solubility

solubility is the quantity of a dissolved substance that is required to form a saturated solution, while Ksp is the equilibrium constant that details the degree of a solid’s dissolving ability

Solubility units (2)

mol/L

g/L

Why does adding a different salt with a common ion to a saturated solution DECREASE the solubility of that original salt?

Since the new salt isn’t yet saturated, the common ion joins with the other ion of the OG salt to form more solid, thus re-establishing equilibrium

Qsp in relation to common ion

At equilibrium, Qsp = Ksp, but adding more common ion increases Qsp

How is equilibrium re-established when adding a small amount of water to a saturated solution?

Since Ksp stays the same and solely equals the concentrations of ions (raised to respective coefficients), the concentrations must stay the same, meaning more moles of solid dissolve to compensate for the greater volume

What if a very large amount of water is added to a saturated solution?

if enough water is added, then the solution will no longer be saturated and equilibrium won’t exist

What happens when adding a strong acid/base to a saturated solution?

As the strong acid/base dissolves into H+ and OH-, the solid’s ions might be strong enough to react further with the H+ or OH-

What should you note about the concentrations of OH- and H3O+ in pure water?

Regardless of temperature, the concentrations of OH- and H3O+ are the same, meaning pH = pOH