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Questions and Answers
Consider the reversible reaction: $2A(g) + B(g) \rightleftharpoons C(g)$. If, at a certain temperature, the equilibrium constant $K_c$ is found to be 4, what does this indicate about the relationship between the rate constants of the forward ($k_f$) and reverse ($k_r$) reactions?
Consider the reversible reaction: $2A(g) + B(g) \rightleftharpoons C(g)$. If, at a certain temperature, the equilibrium constant $K_c$ is found to be 4, what does this indicate about the relationship between the rate constants of the forward ($k_f$) and reverse ($k_r$) reactions?
- The forward rate constant is four times larger than the reverse rate constant ($k_f = 4k_r$). (correct)
- The forward and reverse rate constants are equal ($k_f = k_r$).
- The forward rate constant is four times smaller than the reverse rate constant ($k_f = 4k_r$).
- The forward rate constant is one-fourth the reverse rate constant ($k_f = \frac{1}{4}k_r$).
For the following reaction at equilibrium: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, how will increasing the pressure of the system affect the equilibrium, assuming constant temperature?
For the following reaction at equilibrium: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, how will increasing the pressure of the system affect the equilibrium, assuming constant temperature?
- The equilibrium will randomly fluctuate without a predictable shift.
- The equilibrium will not be affected by the change in pressure.
- The equilibrium will shift to the left, favoring the reactants.
- The equilibrium will shift to the right, favoring the products. (correct)
Consider a reaction where $\Delta n = -1$. This means that Kp will be:
Consider a reaction where $\Delta n = -1$. This means that Kp will be:
- Equal to Kc.
- Less than Kc. (correct)
- Greater than Kc.
- Negative and equal to -Kc.
Which of the following actions will always change the value of the equilibrium constant (Kc) for a given reaction?
Which of the following actions will always change the value of the equilibrium constant (Kc) for a given reaction?
For the reaction $A(aq) + B(aq) \rightleftharpoons 2C(aq)$, the equilibrium constant $K_c$ is 4. If a reaction mixture initially contains [A] = 0.5 M, [B] = 0.5 M, and [C] = 0.0 M, which direction will the reaction proceed to reach equilibrium?
For the reaction $A(aq) + B(aq) \rightleftharpoons 2C(aq)$, the equilibrium constant $K_c$ is 4. If a reaction mixture initially contains [A] = 0.5 M, [B] = 0.5 M, and [C] = 0.0 M, which direction will the reaction proceed to reach equilibrium?
What does a Gibbs Free Energy value of zero indicate about a reaction at a given temperature and pressure?
What does a Gibbs Free Energy value of zero indicate about a reaction at a given temperature and pressure?
In a solution of acetic acid ($CH_3COOH$), the addition of sodium acetate ($CH_3COONa$) will:
In a solution of acetic acid ($CH_3COOH$), the addition of sodium acetate ($CH_3COONa$) will:
Which of the following statements correctly describes the behavior of a buffer solution upon the addition of a strong acid?
Which of the following statements correctly describes the behavior of a buffer solution upon the addition of a strong acid?
What is the key difference between a strong electrolyte and a weak electrolyte in an aqueous solution?
What is the key difference between a strong electrolyte and a weak electrolyte in an aqueous solution?
For a sparingly soluble salt, $AgCl$, what is the correct expression for its solubility product constant, $K_{sp}$?
For a sparingly soluble salt, $AgCl$, what is the correct expression for its solubility product constant, $K_{sp}$?
Flashcards
Equilibrium
Equilibrium
A state where system properties remain constant over time.
Chemical Equilibrium
Chemical Equilibrium
Rate of forward and backward reactions are equal.
Equilibrium Constant (Kc)
Equilibrium Constant (Kc)
Ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients.
Kp
Kp
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Reaction Quotient (Qc)
Reaction Quotient (Qc)
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Gibbs Free Energy (ΔG)
Gibbs Free Energy (ΔG)
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Le Chatelier's Principle
Le Chatelier's Principle
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Electrolytes
Electrolytes
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Brønsted-Lowry theory
Brønsted-Lowry theory
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Buffer solutions
Buffer solutions
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Study Notes
Equilibrium Fundamentals
- Equilibrium is when system properties like temperature, pressure, and concentration are constant.
- Dynamic equilibrium includes continuous motion, like water entering and exiting a bucket at the same rate.
- Chemical equilibrium is achieved when forward and backward reaction rates are equal.
Types of Equilibrium
- Physical equilibrium involves changes to the physical properties of a substance.
- Solid-liquid phase: melting rate equals freezing rate.
- Liquid-vapor phase: evaporation rate equals condensation rate.
- Solid-vapor phase: sublimation rate equals condensation rate.
- Solid-solution phase: dissolution rate equals crystallization rate, exemplified by sugar dissolving in water.
- Chemical equilibrium:
- Homogenous equilibrium: all reactants and products are in the same phase.
- Heterogenous equilibrium: reactants and products are in different phases.
- Reversible reactions reach equilibrium, but irreversible reactions do not.
Graphical Representation of Equilibrium
- Initially, reactant concentration is high, while product concentration is low.
- Over time, reactants decrease as they convert to products.
- Eventually, a dynamic equilibrium is established.
- The Haber process combines N2 and H2 gases to produce ammonia (NH₃).
- H2 and I2 react to form 2HI and reach equilibrium.
Equilibrium Constant (Kc)
- It is the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients.
- Kc equals the forward reaction's constant divided by the backward reaction's constant.
- For the reaction aA + bB ⇌ cC + dD, Kc = [C]^c [D]^d / [A]^a [B]^b, where lowercase letters are the stoichiometric coefficients.
Kp and its Relation to Kc
- Kp is the equilibrium constant expressed using partial pressures.
- Kp = Kc(RT)^Δn
- R is the universal gas constant, T is temperature in Kelvin, and Δn is the change in the number of moles of gas (product moles minus reactant moles).
- If Δn = 0, Kp = Kc; if Δn < 0, Kp < Kc; if Δn > 0, Kp > Kc.
Equilibrium Characteristics and Applications
- Initial reaction concentration does not affect the equilibrium constant.
- Reversing a reaction inverts the equilibrium constant value.
- Multiplying or dividing a reaction by a factor raises the equilibrium constant to that power or root, respectively.
- Kc has different units based on the molar concentrations in the reaction.
- The reaction quotient (Qc) predicts the direction in which a reaction will shift to reach equilibrium.
- At equilibrium, Qc = Kc; Qc < Kc favors the forward reaction; Qc > Kc favors the reverse reaction.
Gibbs Free Energy and Equilibrium
- Gibbs free energy (ΔG) indicates reaction spontaneity.
- ΔG° = -RT ln K_equilibrium or ΔG° = -2.303RT log K_equilibrium
- If K > 1 (ΔG < 0), the reaction is spontaneous and proceeds forward.
- If K < 1 (ΔG > 0), the reaction is non-spontaneous and proceeds backward.
- If K = 1 (ΔG = 0), the reaction is at equilibrium and does not shift.
Le Chatelier's Principle
- Changing factors like temperature, concentration, or pressure shifts the equilibrium to counteract the change.
- Increasing product concentration shifts the reaction backward; increasing pressure shifts the reaction towards fewer gaseous moles.
- Catalysts increase the reaction rate without altering the equilibrium.
Ionic Equilibrium
Electrolytes Overview
- Electrolytes dissociate into ions.
- Strong electrolytes completely dissociate, their degree of dissociation (α) is approximately one, and no equilibrium is established
- Weak electrolytes partially dissociate; their α is less than one, where an ionic equilibrium exists.
- Degree of dissociation measures the fraction of solute molecules that dissociate at a given temperature.
Acid-Base Theories
- Arrhenius theory: acids produce H+ ions, and bases produce OH- ions.
- Brønsted-Lowry theory: acids donate protons (H+), and bases accept protons.
- Lewis theory: acids accept electron pairs, and bases donate electron pairs.
Self-ionization of Water
- Water acts as both an acid and a base (amphoteric), allowing self-ionization.
- Kw = [H+][OH-] = 1.0 x 10^-14 at 25°C.
- Kw changes with temperature.
- pKw = pH + pOH = 14
pH and pOH Concepts
- pH measures the acidity or basicity of a solution.
- pH = -log[H+]; pOH = -log[OH-]
- The pH scale ranges from 0 to 14, with 7 being neutral; values below 7 are acidic, and values above are basic.
- Water's pH is 7 (neutral).
Polyprotic Acids Concepts
- Polyprotic acids contain more than one ionizable hydrogen atom.
- H2SO4 dissociates in two steps, releasing one hydrogen ion (H+) at a time, where the acid dissociation constant changes in each step.
Common Ion Effect Concepts
- Introducing a common ion to a solution at equilibrium shifts the equilibrium.
- Adding NH4Cl to a solution of NH₃ in water increases [NH4+], shifting the reaction backward.
Buffer Solutions Concepts
- Buffer solutions resist pH changes upon the addition of small amounts of acid or base.
- Acidic buffers, basic buffers, and salt buffers are the three main types.
Solubility Concepts
- Solubility measures the extent to which a substance dissolves.
- Substances are soluble if > 0.1M
- Slightly soluble if Between 0.01M and 0.1M
- Sparingly soluble if < 0.0M
- The solubility constant (K_sp) defines the ion product in a saturated solution.
- For a salt A_x B_y, Ksp = [A^+y]^x [B^-x]^y
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