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Equilibrium Flashcards

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1
Q

What is an irreversible reaction?

A

A reaction that takes place in only one direction.

Example: Reaction between NaOH and HCl.

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2
Q

What is a reversible reaction?

A

A reaction that takes place in both directions.

Example: Haber process for the preparation of ammonia.

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3
Q

What is a forward reaction?

A

The process by which reactants are converted to products.

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4
Q

What is a backward reaction?

A

The process by which products recombine to form reactants.

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5
Q

What is equilibrium?

A

A state in which the rates of forward and backward reactions are equal.

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6
Q

Is equilibrium static or dynamic?

A

Dynamic in nature.

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7
Q

What is physical equilibrium?

A

Equilibrium involving physical processes.

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8
Q

What is solid-liquid equilibrium?

A

Equilibrium between ice and water at 273K and 1 atm pressure.

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9
Q

What is liquid-vapour equilibrium?

A

Equilibrium between water and water vapour at 100°C in a closed vessel.

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10
Q

What is solid-vapour equilibrium?

A

Equilibrium established by placing solid iodine in a closed vessel.

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11
Q

What happens during the dissolution of solid in liquids at equilibrium?

A

The rate of dissolution becomes equal to the rate of crystallisation.

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12
Q

What is Henry’s law?

A

The mass of a gas dissolved in a given mass of solvent is proportional to the pressure of the gas above the solvent.

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13
Q

What is chemical equilibrium?

A

Equilibrium associated with chemical reactions, where concentrations of reactants and products are constant.

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14
Q

What does the equilibrium constant (Kc) represent?

A

The ratio of the product of concentrations of products to that of reactants at equilibrium.

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15
Q

What is the relationship between Kc and Kp?

A

Kp = Kc(RT)∆n, where ∆n is the change in number of moles of gaseous species.

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16
Q

What is homogeneous equilibrium?

A

Equilibrium where all reactants and products are in the same phase.

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17
Q

What is heterogeneous equilibrium?

A

Equilibrium where reactants and products are in different phases.

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18
Q

What happens to the equilibrium constant with temperature changes?

A

The value of equilibrium constant depends on temperature.

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19
Q

What does a large value of Kc indicate?

A

The reaction proceeds nearly to completion.

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20
Q

What does the reaction quotient (Qc) represent?

A

The ratio of the concentrations of products to reactants at any point in the reaction.

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21
Q

What does it mean if Qc > Kc?

A

The reaction will proceed in the direction of reactants (reverse direction).

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22
Q

What is Gibb’s energy change related to at equilibrium?

A

ΔG = ΔGΘ + RT lnQc, where ΔGΘ is standard Gibbs energy change.

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23
Q

What principle explains the effect of concentration changes on equilibrium?

A

Le Chatelier’s principle.

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24
Q

How does an increase in concentration of reactants affect equilibrium?

A

Increases the rate of forward reaction, shifting equilibrium to the products side.

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25
What occurs when the concentration of products is increased?
The rate of backward reaction increases, shifting equilibrium to the reactants side.
26
What principle explains the effect of concentration changes on equilibrium?
Le Chatelier’s principle ## Footnote It states that if the concentration of reactants or products is changed, the equilibrium will shift to counteract that change.
27
In the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g), what happens if SO2 or O2 concentration is increased?
The equilibrium shifts to the forward direction ## Footnote This occurs as the system attempts to reduce the concentration of the added reactants.
28
How does temperature affect an exothermic reaction according to Le Chatelier’s principle?
Increase in temperature favors the backward reaction ## Footnote This results in a decrease in product formation.
29
What is the effect of increasing pressure on a gaseous reaction?
It favors the reaction with fewer moles of gas ## Footnote For example, in the reaction N2(g) + 3H2(g) ⇌ 2 NH3(g), an increase in pressure shifts the equilibrium to the right.
30
What role does a catalyst play in an equilibrium reaction?
It increases the rate of both forward and backward reactions equally ## Footnote A catalyst does not change the equilibrium position or composition.
31
What happens when an inert gas is added to an equilibrium mixture at constant volume?
There is no change to the equilibrium ## Footnote The partial pressures or concentrations of the reactants and products remain unchanged.
32
Define ionic equilibrium.
Equilibrium involving ions ## Footnote An example is the dissociation of acetic acid in water.
33
What are electrolytes?
Substances that conduct electricity in molten or solution state ## Footnote Examples include all acids, bases, and most salts.
34
What distinguishes strong electrolytes from weak electrolytes?
Strong electrolytes dissociate almost completely, while weak electrolytes dissociate partially ## Footnote Strong examples: HCl, NaOH; Weak examples: CH3COOH, NH4OH.
35
According to the Arrhenius concept, what defines an acid?
A substance that gives hydrogen ions (H+) in aqueous solution ## Footnote Example: HCl produces H3O+ in water.
36
What does the Bronsted-Lowry concept define as an acid?
Proton (H+) donor ## Footnote Example: NH3 + H2O ⇌ NH4+ + OH−, where NH3 accepts H+.
37
What are Lewis acids and bases defined by?
Lewis acids are electron pair acceptors; Lewis bases are electron pair donors ## Footnote Example: BF3 is a Lewis acid; NH3 is a Lewis base.
38
What is the ionization constant of water (Kw)?
Kw = [H+][OH−] ## Footnote At 298K, Kw = 10^-14 M².
39
What is the pH scale range?
0 to 14 ## Footnote A pH less than 7 indicates acidity, greater than 7 indicates basicity, and equal to 7 indicates neutrality.
40
What is the relationship between pH and pOH?
pH + pOH = 14 ## Footnote This relationship helps in determining the acidity or basicity of a solution.
41
What does Ka represent in acid-base chemistry?
The ionization constant of weak acids ## Footnote It indicates the strength of an acid; larger Ka means a stronger acid.
42
What is the relation between Ka, Kb, and Kw?
Ka x Kb = Kw ## Footnote This applies to conjugate acid-base pairs.
43
How does bond strength affect acid strength?
Weaker H-A bonds lead to stronger acids ## Footnote Acidity increases down a group due to decreasing bond strength.
44
What determines the strength of an acid?
The polarity of the H-A bond and the bond strength ## Footnote Weaker H-A bonds dissociate more easily to give H+ ions, indicating stronger acids.
45
How does acidity change down a group in the periodic table?
Acidity increases as the size of A increases, leading to weaker H-A bonds ## Footnote Example: Acidity of hydrohalic acids increases in the order: HF < HCl < HBr < HI.
46
How does acidity change across a period in the periodic table?
Acidity increases with increasing electronegativity of A ## Footnote Example: Acidity of hydrides in the second period increases in the order: CH4 < NH3 < H2O < HF.
47
What is salt hydrolysis?
The interaction of anion or cation or both of a salt with water ## Footnote This process can affect the pH of the solution.
48
What is the pH of solutions formed from strong acids and strong bases?
pH = 7 (neutral) ## Footnote Cations of strong bases and anions of strong acids do not hydrolyse.
49
What happens during the hydrolysis of a salt formed from a strong base and a weak acid?
Only the anion of the weak acid undergoes hydrolysis, making the solution basic ## Footnote Example: Sodium acetate (CH3COONa) results in pH > 7.
50
What is the formula for calculating the pH of a salt solution from a strong base and weak acid?
pH = 7 + ½ (pKa + log C) ## Footnote C is the concentration of salt.
51
What occurs during the hydrolysis of a salt formed from a weak base and a strong acid?
Only the cation of the weak base undergoes hydrolysis, making the solution acidic ## Footnote Example: Ammonium chloride (NH4Cl) results in pH < 7.
52
What is the formula for calculating the pH of a solution from a weak base and strong acid?
pH = 7 – ½ (pKb + log C) ## Footnote C is the concentration of salt.
53
What is the hydrolysis of a salt formed from weak acid and weak base?
Both cation and anion undergo hydrolysis, producing weak acid and weak base in solution ## Footnote The solution may be neutral, acidic, or basic depending on relative strengths.
54
What is the Henderson-Hasselbalch equation for an acidic buffer?
pH = pKa + log [A−]/[HA] ## Footnote Where HA is the weak acid and A− is its conjugate base.
55
What is the Henderson-Hasselbalch equation for a basic buffer?
pOH = pKb + log [Salt]/[Base] ## Footnote This can be rearranged to pH = 14 - (pKb + log [Base]).
56
What is solubility equilibrium?
It depends on the lattice enthalpy and solvation enthalpy ## Footnote For a salt to dissolve, solvation enthalpy must be greater than lattice enthalpy.
57
How are salts classified based on their solubility?
* Category I: Soluble (Solubility > 0.1M) * Category II: Slightly Soluble (0.01M < Solubility < 0.1M) * Category III: Sparingly Soluble (Solubility < 0.01M)
58
What does Ksp represent?
The solubility product constant ## Footnote It is defined as the product of the molar concentrations of ions in a saturated solution.
59
What is the relationship between Ksp and Qsp?
At equilibrium, Ksp = Qsp ## Footnote If Ksp > Qsp, dissolution occurs; if Ksp < Qsp, precipitation occurs.
60
What is the pH of a solution if [H+] = 3 x 10-3 M?
pH = 2.523 ## Footnote Calculated using pH = -log[H+].
61
What is the formula to calculate the solubility (S) of CaSO4 given Ksp?
S = √Ksp ## Footnote Example: If Ksp = 9 x 10-6, then S = 3 x 10 –3 M.
62
How do you calculate the pH of a 0.01 M acetic acid solution with a degree of ionization of 0.045?
pH = 3.3468 ## Footnote [H3O+] calculated as cα = 0.01 x 0.045.
63
What is the pH of an acidic buffer containing 0.1 M CH3COOH and 0.5 M CH3COONa?
pH = 6.444 ## Footnote Using Henderson-Hasselbalch equation and pKa = 5.7447.
64
What is the solubility product (Ksp) of Mg(OH)2 if its solubility is 1.5 x 10-4 M?
Ksp = 1.35 x 10-11 ## Footnote Ksp calculated as 4S^3 where S = 1.5 x 10-4.
65
What is the pH of a 0.1 M solution of NaOH?
pH = 13 ## Footnote Calculated as 14 - pOH where pOH = 1.

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