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Questions and Answers
In the reaction MnO$_2$ + 4 HCl → MnCl$_2$ + Cl$_2$ + 2 H$_2$O, manganese is reduced and chlorine is oxidized.
In the reaction MnO$_2$ + 4 HCl → MnCl$_2$ + Cl$_2$ + 2 H$_2$O, manganese is reduced and chlorine is oxidized.
True (A)
During the balancing of the half-reaction MnO$_4^−$ → Mn$_2$O$_3$ in acidic solution, water molecules must only be added to the reactant side.
During the balancing of the half-reaction MnO$_4^−$ → Mn$_2$O$_3$ in acidic solution, water molecules must only be added to the reactant side.
False (B)
Electrolytic cells use an external power source to drive a spontaneous redox reaction.
Electrolytic cells use an external power source to drive a spontaneous redox reaction.
False (B)
Increasing the temperature of a reaction will always decrease its rate.
Increasing the temperature of a reaction will always decrease its rate.
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A catalyst shifts the equilibrium of a reaction towards the products.
A catalyst shifts the equilibrium of a reaction towards the products.
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Adding a common ion to a saturated solution of a slightly soluble salt, such as NaCl, will increase the solubility of salt.
Adding a common ion to a saturated solution of a slightly soluble salt, such as NaCl, will increase the solubility of salt.
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A strong acid will have a large Ka value and a small pKa value.
A strong acid will have a large Ka value and a small pKa value.
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In the reaction 2Cr + 3Pb^{2+} \rightarrow 2Cr^{3+} + 3Pb, the standard cell potential, E^o, is -0.61 V.
In the reaction 2Cr + 3Pb^{2+} \rightarrow 2Cr^{3+} + 3Pb, the standard cell potential, E^o, is -0.61 V.
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A reaction with a standard cell potential, E^o, of -0.14V is considered to be spontaneous.
A reaction with a standard cell potential, E^o, of -0.14V is considered to be spontaneous.
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In the reaction 2H_2O + 2I^- \rightarrow 2OH^- + H_2 + I_2, the solution becomes less basic as the reaction proceeds.
In the reaction 2H_2O + 2I^- \rightarrow 2OH^- + H_2 + I_2, the solution becomes less basic as the reaction proceeds.
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In an electrolytic cell using a chromium anode and a cheap metal cathode for plating, chromium ions, Cr^{3+}, are reduced at the cathode.
In an electrolytic cell using a chromium anode and a cheap metal cathode for plating, chromium ions, Cr^{3+}, are reduced at the cathode.
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The collision theory states that the rate of a reaction is determined only by the number of collisions between reactant molecules.
The collision theory states that the rate of a reaction is determined only by the number of collisions between reactant molecules.
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The rate determining step in a reaction mechanism is the fastest step in the reaction.
The rate determining step in a reaction mechanism is the fastest step in the reaction.
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In the multi step reaction: 2A + B → AB + A, AB +C → ABC, ABC + C → AC_2 +B, and AC + AC → A_2C +C, 'A' acts as a catalyst.
In the multi step reaction: 2A + B → AB + A, AB +C → ABC, ABC + C → AC_2 +B, and AC + AC → A_2C +C, 'A' acts as a catalyst.
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For an exothermic reaction, the potential energy of the products is greater than the potential energy of the reactants.
For an exothermic reaction, the potential energy of the products is greater than the potential energy of the reactants.
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The rate law gives information about the effect of pressure on reaction rate.
The rate law gives information about the effect of pressure on reaction rate.
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An increase in temperature typically decreases the reaction rate of a chemical process.
An increase in temperature typically decreases the reaction rate of a chemical process.
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A catalyst alters the activation energy of a reaction, providing a pathway with a higher activation energy.
A catalyst alters the activation energy of a reaction, providing a pathway with a higher activation energy.
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In the reaction $3A \rightarrow 5B$, if the rate of formation of B is 0.08 M/min then the rate of consumption of A is 0.048 M/min.
In the reaction $3A \rightarrow 5B$, if the rate of formation of B is 0.08 M/min then the rate of consumption of A is 0.048 M/min.
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For a reaction that occurs in a single step, the rate law can always be determined solely from the stoichiometry of the reaction.
For a reaction that occurs in a single step, the rate law can always be determined solely from the stoichiometry of the reaction.
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If a reaction is exothermic, the potential energy of the products is lower than the potential energy of the reactants.
If a reaction is exothermic, the potential energy of the products is lower than the potential energy of the reactants.
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The rate determining step of a reaction mechanism is the fastest step in the sequence.
The rate determining step of a reaction mechanism is the fastest step in the sequence.
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An intermediate in a reaction is a stable product which can be isolated from the reaction mixture at the end of the reaction.
An intermediate in a reaction is a stable product which can be isolated from the reaction mixture at the end of the reaction.
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If the concentration of a reactant is doubled and the reaction rate quadruples this suggests that the reaction rate has a second order dependence on that reactant.
If the concentration of a reactant is doubled and the reaction rate quadruples this suggests that the reaction rate has a second order dependence on that reactant.
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In a multi-step reaction mechanism, the sum of the steps must equal the overall reaction stoichiometry.
In a multi-step reaction mechanism, the sum of the steps must equal the overall reaction stoichiometry.
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The rate constant for a reaction generally decreases with an increase in temperature.
The rate constant for a reaction generally decreases with an increase in temperature.
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In the reaction A + B C + D, with a Kc of 64, if 6.0 moles of each reactant are placed in a 4.0 L chamber, the equilibrium concentration of the products will be greater than 1.4 M.
In the reaction A + B C + D, with a Kc of 64, if 6.0 moles of each reactant are placed in a 4.0 L chamber, the equilibrium concentration of the products will be greater than 1.4 M.
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The solubility of calcium sulfate, CaSO4, in a 0.15 M solution of MgSO4 is approximately 2.57 x 10^-4M.
The solubility of calcium sulfate, CaSO4, in a 0.15 M solution of MgSO4 is approximately 2.57 x 10^-4M.
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If iron(II) iodide, FeI2, has a molar solubility of $7.3 x 10^{-9}$ M, then the concentration of Fe$^{2+}$ is $7.3 x 10^{-9}$ M and I$^-$ is $1.46 x 10^{-8}$ M, and the Ksp is about $7.43 x 10^{-24}$
If iron(II) iodide, FeI2, has a molar solubility of $7.3 x 10^{-9}$ M, then the concentration of Fe$^{2+}$ is $7.3 x 10^{-9}$ M and I$^-$ is $1.46 x 10^{-8}$ M, and the Ksp is about $7.43 x 10^{-24}$
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For the reaction A + heat B, at equilibrium, if more of A than B is present, decreasing A will cause the system to shift towards the reactants to reach equilibrium again.
For the reaction A + heat B, at equilibrium, if more of A than B is present, decreasing A will cause the system to shift towards the reactants to reach equilibrium again.
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In the equilibrium reaction $P_4(g) + 6H_2(g) \leftrightarrow 4PH_3(g)$, adding more $H_2(g)$ will shift the equilibrium to the right, increasing the quantity of $PH_3(g)$ at equilibrium.
In the equilibrium reaction $P_4(g) + 6H_2(g) \leftrightarrow 4PH_3(g)$, adding more $H_2(g)$ will shift the equilibrium to the right, increasing the quantity of $PH_3(g)$ at equilibrium.
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For a reaction with the rate law R = k[A]^1[B]^2, the reaction is first order with respect to A and second order with respect to B.
For a reaction with the rate law R = k[A]^1[B]^2, the reaction is first order with respect to A and second order with respect to B.
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If the rate of formation of NH3 is 3.76 x 10^-4 M/s, then the rate of disappearance of H2 is 5.64 x 10^-4 M/s, based on the provided data.
If the rate of formation of NH3 is 3.76 x 10^-4 M/s, then the rate of disappearance of H2 is 5.64 x 10^-4 M/s, based on the provided data.
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Given the rate law R = k[NO]^2[Cl2]^1, and a rate constant k = 0.152 M^-2s^-1, the overall order of the reaction is 4.
Given the rate law R = k[NO]^2[Cl2]^1, and a rate constant k = 0.152 M^-2s^-1, the overall order of the reaction is 4.
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Given the rate of a reaction with respect to A is expressed as Rate of A = 2/3 - ∆[B] / ∆t, this implies that 2 moles of A are consumed for every 3 moles of B consumed.
Given the rate of a reaction with respect to A is expressed as Rate of A = 2/3 - ∆[B] / ∆t, this implies that 2 moles of A are consumed for every 3 moles of B consumed.
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For the equilibrium 2SO3(g) ↔ 2SO2(g) + O2(g), with [SO3] = 0.0160 M, [SO2] = 0.00560 M, and [O2] = 0.00210 M, the equilibrium constant $K_c$ is calculated simply as $(0.00560 * 0.00210) / 0.0160$.
For the equilibrium 2SO3(g) ↔ 2SO2(g) + O2(g), with [SO3] = 0.0160 M, [SO2] = 0.00560 M, and [O2] = 0.00210 M, the equilibrium constant $K_c$ is calculated simply as $(0.00560 * 0.00210) / 0.0160$.
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In the equilibrium N2(g) + 3H2(g) ↔ 2NH3(g) + 92 kJ, increasing the temperature will shift the equilibrium towards the reactants.
In the equilibrium N2(g) + 3H2(g) ↔ 2NH3(g) + 92 kJ, increasing the temperature will shift the equilibrium towards the reactants.
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In the reaction H2(g) + CO2(g) ↔ H2O(g) + CO(g), adding H2(g) would cause the equilibrium to shift to the left.
In the reaction H2(g) + CO2(g) ↔ H2O(g) + CO(g), adding H2(g) would cause the equilibrium to shift to the left.
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For the reaction 4 NH3(g) + 5 O2(g) ↔ 4 NO(g) + 6 H2O(g) with $K_c= 5.0 x 10^{-19}$, the concentrations at equilibrium will predominantly favor the products.
For the reaction 4 NH3(g) + 5 O2(g) ↔ 4 NO(g) + 6 H2O(g) with $K_c= 5.0 x 10^{-19}$, the concentrations at equilibrium will predominantly favor the products.
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If 6.0 mol of CO and 6.0 mol of H2O are placed in a 2.0 L container, with the reaction CO(g) + H2O(g) ↔ H2(g) + CO2(g), and Kc = 0.80, the equilibrium concentration of H2 will always be less than the equilibrium concentration of CO.
If 6.0 mol of CO and 6.0 mol of H2O are placed in a 2.0 L container, with the reaction CO(g) + H2O(g) ↔ H2(g) + CO2(g), and Kc = 0.80, the equilibrium concentration of H2 will always be less than the equilibrium concentration of CO.
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For a reaction where $K_c$ = 3.63, with the concentrations of reactants and products provided, if the calculated Reaction Quotient Q is greater than $K_c$, the reaction shifts towards the reactants to reach equilibrium.
For a reaction where $K_c$ = 3.63, with the concentrations of reactants and products provided, if the calculated Reaction Quotient Q is greater than $K_c$, the reaction shifts towards the reactants to reach equilibrium.
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Study Notes
Chemistry 40S Exam Review
- This review package provides examples of long-answer questions covering the units studied. It does not include all theory; students should also use their notes and unit booklets. Tests and quizzes are also recommended for preparation.
- Topics Covered: Oxidation-Reduction and Electrochemistry, Chemical Kinetics, Chemical Equilibrium and Ksp, Acids and Bases, and Atomic Structure.
Oxidation-Reduction and Electrochemistry
- Includes oxidation numbers, oxidizing/reducing agents, balancing redox reactions (acidic/basic solutions), electrolytic cells, electroplating, electrochemical cells, and corrosion.
Chemical Kinetics
- Covers reaction rates, factors affecting rates (collision theory), reaction mechanisms, and rate laws (reaction orders). Includes catalyzed and uncatalyzed reactions and mechanisms.
Chemical Equilibrium and Ksp
- Includes mass action expressions, equilibrium constants (kc and kp), percent yield of product, ICE charts, Le Chatelier's principle, interpretations of graphs, and calculation of solubility products (ksp) and molar solubilities.
Acids and Bases
- Compares electrolytes and nonelectrolytes, Arrhenius and Brønsted-Lowry definitions, conjugate acid-base pairs, self-ionization of water (Kw), ionization constants (Ka and Kb), pH/pOH scales, calculations, and titration procedures.
Atomic Structure
- Includes the quantum mechanical model of the atom, electron configurations, orbital diagrams, calculations for wavelength, frequency, and energy, types of spectra (continuous, bright line), and periodic trends (atomic/ionic radii, ionization energy, electronegativity).
Electrochemistry Answers (Additional Information)
- Balancing Redox Equations: Shows examples, including half-reaction method and oxidation number method, for balancing chemical equations involving oxidation and reduction.
- Electrochemical Cells: Includes balanced equations, electrochemical cell diagrams, anode/cathode identification, electron flow directions, and calculation of the voltage (E°).
- Spontaneity Prediction: Explains how to determine if a reaction will occur spontaneously and writes out the products of a reaction given the reactants.
- Electrolysis: Given aqueous Nal electrolysis, equations for oxidation and reduction are written. The minimum voltage necessary for the reaction to proceed, and pH changes during the electrolysis process are evaluated.
- Electrolytic Cells and Electroplating: Covers how to construct an electrolytic cell for electroplating, gives the two half-reactions involved and labels the anode, cathode, and electrolyte.
Kinetics (Additional Information)
- Reaction Rates: Includes the factors affecting the rate of a chemical reaction, collision theory, and effective collisions.
- Reaction Mechanisms: Covers reaction mechanisms, rate-determining steps, derivation of possible reaction mechanisms, identification of catalysts and intermediates.
- Potential Energy Graphs: Explains how to draw potential energy diagrams for exothermic reactions, show reactants, products, activated complex, activation energies, and heats of reaction on the graphs, and identify these features when a catalyst is included,
- Calculation of Reaction Rates: Shows how to find the rate of reaction given concentration changes over time, using the example of the consumption of reactants and the formation of products for N2 + 3H2 → 2NH3
- Rate Laws: Includes the relationship of reaction rate to concentrations of reactants and the rate constant.
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Description
Test your knowledge on redox reactions, electrochemistry, and the principles governing these topics. This quiz covers various aspects such as half-reactions, standard cell potentials, and the effects of temperature and catalysts in chemical reactions. Challenge yourself to see how well you understand these fundamental concepts in chemistry.
