Chem Exam 4
What is the pH of a 0.0730 M solution of methylamine (CH₃NH₂, Kb = 4.4 × 10⁻⁴)?
11.75
Determine the pOH of a 0.200 M solution of a weak acid, HA, that is 3.0% dissociated at equilibrium.
11.78
The pH of a 0.0061 M solution of Ca(OH)₂ is
12.09
The pH of a 0.0062 M solution of Ca(OH)₂ is
12.09
If Kp = 573 for the reaction below at 310 K, then what is the value of Kc? 2 A (g) + B (s) ⇌ 2 C (s) + D (g)
14583
What is S° for B in the reaction 3 A → 2 B if ∆S°(rxn) =-163.1 J/mol ･ K? [S° (A) = (205.0 J/mol ･K)]
226.0 J/mol ･ K
What is S° for B in the reaction 3 A → 2 B if ∆S°(rxn) =-161.0 J/mol ･ K? [S° (A) = (205.0 J/mol ･K)]
227.0 J/mol ･ K
What is Ka for the conjugate acid of H₂NNH₂ (Kb = 1.3 × 10⁻⁶)?
7.7 × 10⁻⁹
What is ∆G for a reaction where ∆G° = -4.5 kJ/mol and Q = 3.5 at 295 K? (R = 8.314 J/mol ･ K)
-1.4 kJ/mol
Consider the equilibrium system described by the chemical reaction below, which has a value of Kp equal to 4.7 × 10⁻¹⁰ at 200 K. If a flask is initially charged with 2.00 atm of H₂S and 5.00 atm of O₂, what will the equilibrium partial pressure of SO₂ be?
0.0218 atm
For the galvanic (voltaic) cell Fe(s) + Mn²⁺(aq) ⟶ Fe²⁺(aq) + Mn(s) (E° = 0.77 V at 25°C), what is [Fe²⁺] if [Mn²⁺] = 0.060 M and E = 0.78 V? Assume T is 298 K
0.028 M
Consider the equilibrium system described by the chemical reaction below, which has a value of Kc equal to 1.87 × 10⁻³ at a certain temperature. If 5.00 g of solid PH₃BCl₃ and 0.0700 g of BCl₃ are added to a 4.500 L reaction vessel, what will the equilibrium concentration of PH₃ be?
0.0432 M
Which of the followings statements is (are) true for a galvanic (voltaic) cell when E° = +1.00 V? 1. The reaction is spontaneous 2. At equilibrium, K = 1. 3. ∆G° is negative.
1 and 3 only
For the reaction CCl₄(g) ⇌ C(s) + 2Cl₂(g) Kp = 0.76 at 700°C. What initial pressure of pure carbon tetrachloride will produce a TOTAL equilibrium pressure of 1.60 at 700°C?
1.21 atm
Determine the mass of solid NaCH₃COO that must be dissolved in an existing 500.0 mL solution of 0.200 M CH₃COOH to form a buffer with a pH equal to 5.00. The value of Ka for CH₃COOH is 1.8 × 10⁻⁵.
15g
A light bulb consumes energy at a rate of 80 joules per second. How long in seconds will it take for the light bulb to consume 2.30 × 10^5 J in energy?
2880 s
What is the value of n in the Nernst equation for the reaction Al(s) + 3 Ag⁺(aq) ⟶ Ag(s) + Al³⁺(aq).
3
What is the pH of a solution of 0.300 M HNO₂ containing 0.230 M NaNO₂? (Ka of HNO₂ is 4.5 × 10⁻⁴) Correct
3.23
Carbonic acid, H₂CO₃ is a diprotic acid with Ka1 = 4.3 × 10⁻⁷ and Ka2 = 5.6 × 10⁻¹¹. What is the pH of a 0.60 M solution of carbonic acid?
3.29
Consider the equilibrium system described by the chemical reaction below. A 1.00 L reaction vessel was filled with 2.00 mol SO₂ and 2.00 mol NO₂ and allowed to react at a high temperature. At equilibrium, there were 1.30 mol of NO in the vessel. Determine the concentrations of all reactants and products at equilibrium and then calculate the value of Kc for this reaction.
3.4
What is the solubility of Cr(OH)₃ at a pH of 9.10? (Ksp Cr(OH)₃ is 6.70 × 10⁻³¹)
3.4 × 10⁻¹⁶ M
What is the pH of a 1.0 L buffer made with 0.300 mol of HF (Ka = 6.8 × 10⁻⁴) and 0.200 mol of NaF to which 0.150 mol of NaOH were added?
3.54
The pH of a solution is 3.60. What is the OH⁻ concentration in the solution?
4.0 × 10⁻¹¹ M
Determine the pH of a buffer formed by dissolving 21.5 g HC₇H₅O₂ and 37.7 g of NaC₇H₅O₂ in 200.0 mL of solution. The value of Ka for HC₇H₅O₂ is 6.3 × 10⁻⁵.
4.37
You would like to produce a gold-plated coin by plating gold onto a penny 1.90 cm in diameter. How many days will it take to produce a layer of gold 0.480 mm thick (on both sides of the coin) from an Au³⁺ bath using a current of 0.0200 A? (density of gold = 19.3 g/cm³) For the purposes of this problem, you can ignore the edges of the coin.
4.47 days
Determine the pH at the point in the titration of 40.0 mL of 0.200 M H₂NNH₂ with 0.100 M HNO₃ after 80.0 mL of the strong acid has been added. The value of Kb for H₂NNH₂ is 3.0 × 10⁻⁶.
4.83
Determine ∆S for the phase change of 1.85 moles of water from solid to liquid at 0°C. (∆H = 6.01 kJ/mol)
40.7 J/K
The pOH of a solution is 9.87. What is the H⁺ concentration in the solution?
7.4 × 10⁻⁵ M
Determine the resulting pH when 0.040 mol of solid NaOH is added to a 200.0 mL buffer containing 0.100 mol C₆H₅NH₃Cl and 0.500 M C₆H₅NH₂. The value of Kb for C₆H₅NH₂ is 4.3 × 10⁻¹⁰.
5.00
You want to quickly set up a temporary water bath in your lab with a volume of 10.0 L and a temperature of 37.0°C. You only have hot water from your hot water faucet (temperature = 51.8°C) and cold water from your cold water faucet (temperature = 22.0°C). What volume of hot water (in liters) must you mix with cold water to get 10.0 L of 37.0°C water? Assume the specific heat of the water is 4.184 J/g･K and that the water has a density of 1.00 g/mL.
5.03 L
What is the minimum mass of Mg(NO₃)₂ that must be added to 1.00 L of a 0.310 M HF solution to begin precipitation of MgF₂(s)? For MgF₂, Ksp = 7.4 × 10⁻⁹, and Ka for HF = 7.2 × 10⁻⁴.
5.16 × 10⁻³ g
An unknown weak base with a concentration of 0.0910 M has a pH of 12.30. What is the Kb of this base?
5.6 × 10⁻³
A 122.3 g piece of copper (specific heat 0.38 J/g･°C) is heated and then placed into 400.0 g of water initially at 20.7°C. The water increases in temperature to 22.2°C. What is the initial temperature of the copper? (The specific heat of water is 4.18 J/g･°C)
76.2 °C
Consider the following equilibrium reaction: A (g) + B (s) ⇌ C (g) If Kp = 6 × 10⁻³, which species will have the highest partial pressure at equilibrium?
A
Why does aluminum (Al) metal not undergo corrosion like iron does?
A protective layer of Al₂O₃ forms on the surface of Al.
Which of the following is the cell diagram for the reaction 3 Pb²⁺+(aq) + 2 Al(s) → 3 Pb(s) + 2 Al³⁺(aq)?
Al(s) | Al³⁺(aq) || Pb²⁺(aq) | Pb(s)
What is the conjugate base of the Brønsted-Lowry acid, CH₃CH₂COOH?
CH₃CH₂COO⁻
Consider the reaction below. Which species is(are) the Brønsted-Lowry base(s)? HCO₃⁻ (aq) + F⁻ (aq) ⇌ CO₃²⁻ (aq) + HF (aq)
CO₃²⁻, F⁻
Which pairs of species will spontaneously react?
Cd and Sn²⁺
In the electrochemical cell Ni(s) | Ni²⁺(1 M) || H⁺(1 M) | H₂(1 atm) | Pt(s), which change will cause E of the cell to decrease?
Double the pressure of H₂
Which metal can be a sacrificial anode for an iron pipe?
Cr
A galvanic cell is represented by the shorthand Cu | Cu²⁺ || Ag⁺ | Ag. Which reaction occurs at the anode?
Cu(s) → Cu²⁺(aq) + 2e⁻
Which of the following aqueous solutions will not be more soluble under acidic conditions?
CuBr
Which of the following has the smallest standard molar entropy (∆S°)?
Diamond
For the electrochemical cell represented by Cr²⁺ | Cr³⁺ || Li⁺ | Li, E°(red) = -2.63 V. This cell:
Does not run spontaneously and is electrolytic
Consider the following half-reactions: I₂(s) + 2e⁻ → 2 I⁻(aq), E°(red) = 0.54 V, Cd²⁺(aq) + 2 e⁻ → Cd(s), E°(red) = -0.40 V. Will the electrochemical cell represented by I⁻ | I₂ || Cd²⁺ | Cd be galvanic or electrolytic under standard conditions?
Electrolytic
Evaluate the validity of the following statement: Spontaneous processes are ones that occur quickly and have a low activation energy.
False. Spontaneous reactions can react slowly and can have a high activation energy.
The correct reaction showing how FeCO₃ has increased solubility when forming the complex ion Fe(CN)₆⁴⁻ is
FeCO₃ (s) + 6 CN⁻ (aq) ⇌ Fe(CN)₆⁴⁻ (aq) + CO₃²⁻ (aq)
In the redox reaction 6 Fe²⁺ + Cr₂O₇²⁻ + 14 H⁺ → 2 Cr³⁺ + 6 Fe³⁺ + 7 H₂O, what is the reducing agent?
Fe²⁺
Which one of the following statements is false?
For a triprotic acid, Ka3 > Ka2
Rank the following solutions in order from most acidic to least acidic.
HNO₃ > HNO₂ > NaNO₃ > NaNO₂
Which of the following equations represents the acid equilibrium associated with Ka₃ for H₃PO₃?
HPO₃²⁻ (aq) + H₂O (l) ⇌ PO₃³⁻ (aq) + H₃O⁺ (aq)
Which of the following is an amphoteric species?
HSO₃⁻
Consider the reaction below. Which species are conjugate acid/base pairs? HSO₃⁻ (aq) + HCN (aq) ⇌ H₂SO₃ (aq) + CN⁻ (aq)
HSO₃⁻, H₂SO₃
What is the effect on the concentration of H⁺ in a solution of HCN if KCN is added to the equilibrium solution?
It decreases
Which of the following statements concerning equilibrium is not true?
The equilibrium constant is independent of temperature
Which of the following species is the best reducing agent?
Mg
Which of the following is a means of creating a buffer of H₂CO₃/NaHCO₃?
Mixing 5 mL of 1 M HNO₃ with 10 mL of 1 M NaHCO₃
What is the primary species in solution at the halfway point in a titration of NH₃ with HBr?
NH₃ and NH₄⁺
Which of the following reactions would have the most negative ∆S° value ?
NH₃(g) + HBr(g) ⟶ NH₄Br(s)
A 0.600 M solution of NaX has a pH of 11.19. Which of the following could be NaX?
NaBrO
What is the conjugate acid of the Brønsted-Lowry base, N₂H₄?
N₂H₅⁺
The conjugate base of HPO₃²⁻ is
PO₃³⁻
Consider the reaction in a lead storage battery below. Which of the following statements is correct? PbO₂(s) + Pb(s) + 2 H₂SO₄(aq) ⇌ 2 PbSO₄(s) + 2 H₂O(l)
Pb²⁺ is formed at the cathode during use of the battery.
What does a 'dead' battery mean chemically?
Q = K
Which is true for the endothermic reaction 2 SO₃(g) → 2 SO₂(g) + O₂(g)?
Spontaneous at high T
If E°(red) of a given half-cell is more negative than E°(red) for a standard hydrogen electrode, the half-cell will:
Tend to oxidize
An increase of pH by 2 implies
The H⁺ concentration decreases by a factor of 100.
A solution has a pH of 13.20. We can conclude that
The solution is very basic
Rank the following in order of most acidic to least acidic.
Z > X > Y
A galvanic cell Zn | Zn²⁺ || Ni²⁺ | Ni runs spontaneously. If a current is imposed to turn this into an electrolytic cell, which of the following will occur?
Zn²⁺ gets reduced
It is found that up to 0.0110 g of SrF₂ dissolves in 100 mL of aqueous solution at a certain temperature. Determine the value of Ksp for SrF₂.
[0.0110]^2
A metal oxide will form _______ and a nonmetal oxide will form _________ in aqueous solution.
a base, an acid
If the Gibbs free energy for an equilibrium is a large, negative number, the equilibrium constant is expected to be _____
a large, positive value
Which of the following would decrease the solubility of a 0.10 M solution of Ag₂CO₃ the most? (Ksp of Ag₂CO₃ is 8.1 × 10⁻¹²)
adding 0.10 M Ag
When NO₂ dimerizes (two molecules join together to form a 'dimer') into N₂O₄, an equilibrium is reached and in the process this reaction produces heat. If you increase the temperature of the chamber in which both NO₂ and N₂O₄ reside, you would observe
an increase in NO₂
In a refrigerator's expansion valve, the pressure on the refrigerant _____ as it vaporizes, an _____ process.
decreases; endothermic
The value of Kw increases as temperature increases. Is the autoionization of water exothermic or endothermic?
endothermic
What type of property is heat?
extensive physical property
At the equivalence point for the titration of HCN with KOH, the pH is expected to be
greater than 7
In examples where the volume of a system is under constant pressure, which of the following quantities equals enthalpy change?
heat absorbed or released by the system
The solubility of a salt refers to
how much of a salt will dissolve
PbCl₂ would have the highest solubility in
pure water
A student wishes to determine the chloride ion concentration in a water sample at 25 °C using a galvanic cell constructed with a graphite electrode and a half-cell of AgCl(s) + e⁻ → Ag(s) + Cl⁻(aq) E°red = 0.2223 V And a copper electrode with 0.500 M Cu²⁺ as the second half cell Cu²⁺(aq) + 2 e⁻ → Cu(s) E°red= 0.337 V The measured cell potential when the water sample was placed into the silver side of the cell was 0.0829 V.
on my phone
An equilibrium is established for the endothermic reaction 2 CO(g) + MoO₂(s) ⇌ 2 CO₂(g) + Mo(s). How would each of the following changes affect the partial pressure of carbon dioxide at equilibrium?
on my phone
Which of the following measures the average kinetic energy of a system?
temperature
Electrons always flow in a voltaic (galvanic) cell from _____
the anode to cathode
The reason complex ion formation can increase the solubility of insoluble compounds is
the complex ions shift the solubility equilibrium forward by removing the cation from the solution
Work is best described as
the displacement of an object against a force
A solution has a pOH of 7.84. This tells us
the solution is only slightly acidic.
Which of the following does the change in the free energy of a reaction predict?
the spontaneity
Determine if the following salt is neutral, acidic or basic. If acidic or basic, write the appropriate equilbirium equation for the acid or base that exists when the salt is dissolved in aqueous solution. If neutral, write only NR. (CH₃)₃NHBr
(CH₃)₃NH⁺(aq) + H₂O(l) ⇌ (CH₃)₃N(aq) + H₃O⁺(aq)
Consider the following reaction at 25 °C: 3 NiO(s) + 2 NH₃(g) → 3 Ni(s) + N₂(g) + 3 H₂O(g)
(answers and explanations on phone)
For the gas-phase equilibrium A(g) + 2 B(g) ⇌ C(g) the initial partial pressures of A, B, and C are all 0.300 atm. After equilibrium is established at 25°C, it is found that the partial pressure of C is 0.190 atm. What is ∆G° for this reaction? (R = 8.314 J/mol･K).
-1.33 kJ/mol
At 1120 K, ∆G° = 89.5 kJ/mol for the reaction 3 A (g) + B (g) →2 C (g). If the partial pressures of A, B, and C are 11.5 atm, 8.60 atm, and 0.510 atm respectively, what is the free energy for this reaction?
-11.3 kJ/mol
Determine ∆G° for a reaction when ∆G = -179.3 kJ/mol and Q = 0.043 at 298 K. (R = 8.314 J/mol ･ K)
-171.5 kJ/mol
What is the maximum amount of work that is possible for an electrochemical cell where E = 1.14 V and n = 2? (F = 96,500 J/(V･mol))
-2.2 × 10⁵ J
Consider the following half reaction: Na⁺(aq) + e⁻ → Na(s). For this reaction, E°(red) = -2.7 V. If this reaction is tripled so that 3 Na⁺ ions are reduced to 3 Na atoms, what is the new E°(red)?
-2.7 V
For the equilibrium A + B ⇌ C + D K was measured to be 1.20 at 25°C and 1.95 × 10⁻² at 300°C. What is ∆H° for the reaction?
-21.3 kJ/mol
Using the provided table and the equation below, determine the heat of formation for KClO₂. 2 KClO₃ (s) → 2 KClO₂ (s) + O₂ (g) ∆H° = 296.2 kJ/mol
-243.1 kJ/mol
A solid 30.2 cm^3 block of KClO3 is heated in the laboratory and decomposes according to the following reaction. What is the change in enthalpy in kJ when all the KClO3 decomposes? The density of KClO3 is 2.34 g/cm^3.
-25.8 kJ
Consider a balloon filled with 2115 L of helium at 1.00 atm pressure and 45.5°C. The temperature of the balloon is decreased to 19.1°C with the pressure remaining at 1.00 atm. What is ∆E (in kJ) for this change? (Assume ideal conditions. The molar heat capacity for He at constant pressure, Cp, is 20.8 J/°C･mol.)
-26.7 kJ
How much heat will be released (∆H) if 0.4247 mol of NH₃ are mixed with 0.20 mol of O₂ in the following chemical reaction? 4 NH₃ (g) + O₂ (g) → 2 N₂H₄ (g) + 2 H₂O (g) ∆H° = -286 kJ/mol
-30.4 kJ
7.71 g of MgSO₄ is placed into 100.0 mL of water. The water's temperature increases by 6.70°C. Calculate ∆H, in kJ/mol, for the dissolution of MgSO₄. (The specific heat of water is 4.18 J/g･°C and the density of the water is 1.00 g/mL). You can assume that the specific heat of the solution is the same as that of water.
-47.1 kJ/mol
The neutralization of a strong acid and strong base has an enthalpy change, ∆H°, = -55.9 kJ/mol. The net ionic equation for this reaction is H⁺(aq) + OH⁻(aq) ⇌ H₂O(l) which is the reverse of the autoionization of water, and therefore its equilibrium constant, K = 1.0 × 10¹⁴ at 25°C. You can determine ∆S° from this data. Use the ∆H° and ∆S° values to calculate ∆G° at 80°C. Assume ∆H° and ∆S° are temperature independent.
-84.3 kJ/mol
The pH for 0.0715 M solution of CCl₃CO₂H is 1.40. Determine the value of Ka for CCl₃CO₂H.
0.050
Consider the equilibrium system described by the chemical reaction below. For this reaction, Kc = 2.4 × 10⁻³ at a particular temperature. If the equilibrium concentrations of H₂O and H₂ are 0.11 M and 0.019 M, respectively, determine the concentration of O₂ at equilibrium..
0.080 M
The pH for 0.185 M solution of an unknown weak base, B, is 12.95. Determine the value of Kb for the unknown base.
0.083
Consider the equilibrium system described by the chemical reaction below. When 1.50 mol of CO₂ and an 10.0 mol solid carbon are heated in a 20.0 L container at 1100 K, the equilibrium concentration of CO is 0.0700 M. Determine the concentrations of all species at equilibrium and then calculate the value of Kc for this reaction.
0.123
What is the percent ionization of H₂NNH₂ in a solution with a concentration of a 0.520 M? (Kb = 1.3 × 10⁻⁶)
0.16 %
Consider a solution made by mixing 500.0 mL of 4.0 M NH₃ and 500.0 mL of 0.40 M AgNO₃. Ag⁺ reacts with NH₃ to form AgNH₃⁺ and Ag(NH₃)₂⁺: Ag⁺ + NH₃ <--> AgNH₃⁺ K₁ = 2.1 × 10³ AgNH₃⁺ + NH₃ <--> Ag(NH₃)₂⁺ K₂ = 8.2 × 10³ The concentration of Ag(NH₃)₂⁺ at equilibrium is
0.20 M
For the reaction A (g) → 3 B (g), Kp = 11300 at 298 K. When ∆G = -14.2 kJ/mol, what is the partial pressure of A when the partial pressure of B is 2.00 atm for this reaction at 298 K.
0.2183 atm
Consider the equilibrium system described by the chemical reaction below. Calculate the value of Qc for the initial set reaction conditions in a 2.00 L container: 0.0560 g H₂, 4.36 g I₂, and 1.32 g HI.
0.223
A 20.0 mL solution of NaOH is neutralized with 30.0 mL of 0.200 M HBr. What is the concentration of the original NaOH solution?
0.300 M
Which of the following buffers has the largest capacity if made in 1.0 L solutions?
0.300 mol HCN and 0.200 mol NaCN
Consider the equilibrium system described by the chemical reaction below, which has a value of Kp equal to 60.6 at a certain temperature. If an initial mixture of 0.20 atm of every species is allowed to react, what will the equilibrium partial pressure of NO be?
0.37 atm
An empty steel vessel is charged with 0.250 atm of C and 0.250 atm of A. Once the system reaches equilibrium according to the reaction below, what is the equilibrium partial pressure of C? Kp for this reaction is 11.2. A (g) ⇌ B (g) + C (g)
0.490 atm
Which of the following solutions has the highest H⁺ concentration?
0.50 M HF and 0 M NaF
Determine E° for the half-reaction Fe³⁺(aq) + e⁻ → Fe²⁺(aq).
0.77 V
Which of the following are spontaneous processes? 1. Ice melting at 273 K (assume only ice is initially present). 2. Heat flowing from a hot object to a cold object. 3. An iron bar rusting.
1, 2, and 3
What is the maximum concentration of Ni²⁺ that can be added to a 0.00130 M solution of Na₂CO₃ before a precipitate will form? (Ksp for NiCO₃ is 1.30 × 10⁻⁷)
1.00 × 10⁻⁴ M
The OH⁻ concentration in an aqueous solution at 25 °C is 9.1 × 10⁻³. What is [H⁺]?
1.1 × 10⁻¹² M
What is the entropy change when 329 J of energy is reversibly transferred to a sample of water at 25°C?
1.10 J/K
In an aqueous solution at 25°C, if [H₃O⁺] = 8.1 × 10⁻⁴ M, then [OH⁻] is:
1.2 × 10⁻¹¹ M
A voltaic cell using Cu²⁺/Cu and Al³⁺/Al half-cells is set up at standard conditions, and each compartment has a volume of 225 mL. What is the [Al³⁺] after the cell has delivered 0.120 A for 39.0 hours at 25°C? (E° for Cu²⁺/Cu = 0.340 V and E° for Al³⁺/Al = -1.660 V.)
1.26 M
In a galvanic cell based on the half reactions Ni²⁺(aq) + 2e⁻ → Ni(s) E° = -0.250 V 2H⁺(aq) + 2e⁻ → H₂ E° = 0.000 V the nickel compartment contains a nickel electrode in a solution where [Ni²⁺] = 1.00 × 10⁻² M, and the hydrogen compartment contains a platinum electrode, P_H₂ = 1.00 atm, and a weak acid, HA, at an initial concentration of 1.00 M. If the observed cell potential is 0.135 V at 25°C, calculate the Ka value for the weak acid.
1.29 × 10⁻⁶
A solution of phenol (HC₆H₅O) is prepared by dissolving 0.385 g of phenol in 2.00 L of H₂O. The solution has a pH = 6.29. Determine the value of Ka for phenol.
1.3 × 10^-10
Determine the molar solubility for Cr(OH)₃ (Ksp = 6.3 × 10⁻³¹) in an aqueous solution that has a pH of 11.90 at 25 °C.
1.3x10^-24
Excess solid Na₂CO₃ is added to a solution containing 0.500 M (each) Mg²⁺ and Zn²⁺ ions. Ksp for MgCO₃ is 3.50 × 10⁻⁸ and Ksp for ZnCO₃ is 1.00 × 10⁻¹⁰. ZnCO₃, with the smaller Ksp, will be the least soluble and will begin precipitating first. What will be the [Zn²⁺] concentration when MgCO₃ just begins to precipitate? (Assume no volume change upon addition of the solid Na₂CO₃).
1.43 × 10⁻³ M
For the electrochemical cell 2 Al(s) + 3 Mn²⁺(aq) ⟶ 2 Al³⁺(aq) + 3 Mn(s) (E° = 0.48 V, [Al³⁺] = 1.0 M), which of the following concentrations of Mn²⁺ would cause an increase in E for the cell?
1.5 M
A neutral solution of water at a particular temperature has a concentration of OH⁻ of 4.1 × 10⁻⁷ M. What is Kw at this temperature?
1.7 × 10⁻¹³
What is the solubility of La(IO₃)₃ in a solution that contains 0.350 M IO₃⁻ ions? (Ksp of La(IO₃)₃ is 7.5 × 10⁻¹²).
1.7 × 10⁻¹⁰ M
Determine concentration of OH⁻ in a 0.724 M solution of BrO⁻ (Kb = 4.0 × 10⁻⁶).
1.7x10^-3
How many moles of Fe(OH)₂ [Ksp = 1.8 × 10⁻¹⁵] will dissolve in 1 L of water buffered at pH = 12.00?
1.8 × 10⁻¹¹ mol
An empty steel container is filled with 3.60 atm of H₂ and 3.60 atm of F₂. The system is allowed to reach equilibrium. If Kp = 0.450 for the reaction below, what is the equilibrium partial pressure of HF? H₂ (g) + F₂ (g) ⇌ 2 HF (g)
1.81 atm
The concentration of Mg²⁺ in seawater is 0.052 M. At what pH will 99% of the Mg²⁺ be precipitated as the hydroxide? (Ksp for Mg(OH)₂ = 8.9 × 10⁻¹²)
10.12
What is the pH of a solution of 0.400 M CH₃NH₂ containing 0.230 M CH₃NH₃I? (Kb of CH₃NH₂ is 4.4 × 10⁻⁴)
10.88
At 355 K, ∆G = -5.7 kJ/mol for the reaction A (g) → 2 B (g). If the partial pressures of A and B are 3.37 atm and 0.110 atm respectively, what is the standard free energy for this reaction at this temperature?
10.92 kJ/mol
Which of the following will not produce a buffered solution?
100 mL of 0.1 M Na₂CO₃ and 50 mL of 0.1 M KOH
Calculate ∆E when 2 moles of a liquid is vaporized at its boiling point (85°C) and 1.00 atm pressure. ∆Hvap for the liquid is 55.7 kJ/mol at 85°C. Recall ∆H = ∆E + P∆V and watch your units!
105.4 kJ
Consider the equilibrium system described by the chemical reaction below. At equilibrium, a 6.5 L reaction vessel contained a mixture of 7.38 atm H₂, 2.46 atm N₂, and 0.166 atm NH₃ at 472 °C. What are the values of Kp and Kc for the reaction at this temperature?
109
What is the pH of a 0.420 M solution of NaCN (Ka of HCN is 4.9 × 10⁻¹⁰)?
11.47
What is the pH of a 0.620 M solution of NaCN (Ka of HCN is 4.9 × 10⁻¹⁰)?
11.55
For the following reaction 2NO₂(g) ⇌ N₂O₄(g) ∆H° = -58.02 kJ/mol and ∆S° = -176.6 J/mol･K. What is ∆G for the reaction at 95.0°C when P_NO₂ = P_N₂O₄ = 0.200 atm? Assume ∆H° and ∆S° are temperature independent.
11.89 kJ/mol
For the reaction below the partial pressure of NO₂ is 0.23 atm and the partial pressure of N₂O₄ is 1.0 atm. What is the Q of the reaction? 2 NO₂ (g) ⇌ N₂O₄ (g) Kp = 0.25
19
Which one of the following reactions would produce the largest amount of heat per mole of oxygen?
2 Ca (s) + O₂ (g) → 2 CaO (s) ∆H° = -635 kJ/mol
During the electrolysis of molten NaI, what reaction occurs at the anode?
2 I⁻(l) → I₂(g) + 2e⁻
Complete the balanced molecular reaction for the following weak base with a strong acid: 2 NaClO₂(aq) + H₂SO₄(aq) →
2 NaClO₂(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2 HClO₂(aq)
(no work shown) How many moles of HCl(g) must be added to 1.0 L of 1.0 M NaOH to achieve a pH of 0.00? (Neglect any volume change.)
2.0 mol
The pOH of an acidic solution is 10.63. What is [OH⁻]?
2.3 × 10⁻¹¹ M
Consider the following reactions: A ⇌ B, K₁=4.62 A ⇌ C, K₂=2.00 What is K for the reaction C ⇌ B?
2.31
(doesn't have work)Consider the following equilibrium: N₂(g) + 3H₂(g) <--> 2NH₃(g) with K = 2.3 × 10⁻⁶. 1.00 mol each of all reactants and products is placed in a 1.00-L container. Calculate the equilibrium concentration of H₂.
2.5 M
Aspirin (acetylsalicylic acid, HC₉H₇O₄) has a value of Ka equal to 3.3 × 10⁻⁴. What is the pH after 652 mg of aspirin is dissolved in a solution of 237 mL?
2.68
Consider the equilibrium system described by the chemical reaction below. A 2.00 L reaction vessel was filled 0.0432 mol SO₂ and 0.0296 mol O₂ at 900 K and allow to react. At equilibrium, the concentration of SO₃ was found to be 0.0175 M. Determine the concentrations of all species at equilibrium and then calculate the value of Kc for this reaction.
2.99x10^3
Using the equations 2 Sr(s) + O₂ (g) → 2 SrO (s) ∆H° = -1184 kJ/mol SrO (s) + CO₂ (g) → SrCO₃ (s) ∆H° = -234 kJ/mol CO₂ (g) → C(s) + O₂ (g) ∆H° = 394 kJ/mol Determine the enthalpy for the reaction 2 SrCO₃ (s) → 2 Sr (s) + 2 C(s) + 3 O₂ (g).
2440 kJ/mol
At what temperature, in °C, is a certain reaction at equilibrium if ∆H = +88.8 kJ/mol and ∆S = +170.2 J/mol ･ K?
249 °C
What is the maximum concentration of Ag⁺ that can be added to a 0.00680 M solution of Na₂CO₃ before a precipitate will form? (Ksp for Ag₂CO₃ is 8.10 × 10⁻¹²)
3.45 × 10⁻⁵ M
The solubility of Ni₃(PO₄)₂ in water at a particular temperature is 8.0 × 10⁻⁴ M. What is Ksp for Ni₃(PO₄)₂?
3.5 × 10⁻¹⁴
What is the solubility of Ca(OH)₂ in 0.670 M Ba(OH)₂? Ksp for Ca(OH)₂ is 6.50 × 10⁻⁶. Ba(OH)₂ is a strong base
3.62 × 10⁻⁶ M
What is the pH of a buffer made from 0.200 mol of HCNO (Ka = 3.5 × 10⁻⁴) and 0.410 mol of NaCNO in 2.0 L of solution?
3.77
Using the equations H₂ (g) + F₂ (g) → 2 HF (g) ∆H° = -79.2 kJ/mol C (s) + 2 F₂ (g) → CF₄ (g) ∆H° = 141.3 kJ/mol Determine the enthalpy for the reaction C (s) + 4 HF (g) → CF₄ (g) + 2 H₂ (g).
300 kJ/mol
How many moles of H₂ are required to produce -9301 kJ of heat in the following reaction? N₂ (g) + 3 H₂ (g) → 2 NH₃ (g) ∆H° = -91.8 kJ/mol
304.0 mol
When 0.200 mol of CaCO₃(s) and 0.300 mol of CaO(s) are placed in an evacuated, sealed 10.0-L container and heated to 385 K, PCO₂ = 0.220 atm after equilibrium is established. CaCO₃(s) ⇌ CaO(s) + CO₂(g) Additional CO₂(g) is pumped into the container to raise the pressure to 0.810 atm. After equilibrium is re-established, what is the total mass (in g) of CaCO₃ in the container?
31.7 g
A ball has a kinetic energy of 7.60 kJ. If the ball has a mass of 120.0 g, how fast is the ball traveling?
356 m/s
The Kp for the reaction A (g) ⇌ 2 B (g) is 0.0510. What is Kp for the reaction 4 B (g) ⇌ 2 A (g)?
384
Determine the molar solubility for Cd₃(PO₄)₂ (Ksp = 2.5 × 10⁻³³) in a solution that already contains 0.500 M Na₃PO₄.
4.5 × 10⁻12 M
A cubic piece of platinum metal (specific heat capacity = 0.1256 J/°C･g) at 200.0°C is dropped into 1.00 L of deuterium oxide ('heavy water,' specific heat capacity - 4.211 J/°C･g) at 25.5°C. The final temperature of the platinum and deuterium oxide mixture is 35.3°C. The density of platinum is 21.45 g/cm³ and the density of deuterium oxide is 1.11 g/mL. What is the edge length of the cube of platinum?
4.7 cm
For the reaction: 2 A (g) + B (s) ⇌ 2 C (s) + D (g) Kp = 8210 At 298 K in a 10.0 L vessel, the known equilibrium values are as follows: 0.071 atm of A, 0.22 mol of B, and 10.5 mol of C. What is the equilibrium partial pressure of D?
41.4 atm
The vaporization of methanol CH₃OH(l) ⇌ CH₃OH(g) is an endothermic process in which ∆Hvap = 37.4 kJ/mol and in which entropy increases, ∆Svap = 111 J/mol･K. Calculate the boiling point of methanol (in °C) on a mountaintop where the atmospheric pressure is 305 torr. (Assume ∆H° and ∆S° do not change appreciably with temperature).
42.4 °C
How much heat would be required to convert 133.7 g of ice to water at 0°C? (∆Hfus = 6.01 kJ/mol for water)
44.59 kJ
With which of the following solutions would a 3.5 × 10⁻⁴ M solution of Ag⁺ ions NOT form a precipitate?
5.0 × 10⁻⁵ M CO₃²⁻
The molar solubility of Zn₃(PO₄)₂ is 5.6 × 10⁻⁵ M at a certain temperature. Determine the value of Ksp for Zn₃(PO₄)₂.
5.9x10^-20
How many milliliters of 0.0200 M Ca(OH)₂are required to neutralize 73.2 mL of 0.0300 M HCl?
54.9 mL
What is the equilibrium constant K at 25°C for an electrochemical cell when E° = +0.0810 V and n = 2?(F = 96,500 J/(V･mol), R = 8.314 J/(mol･K))
549.94
A neutral solution of water at a particular temperature has a concentration of OH⁻ of 7.8 × 10⁻⁷ M. What is Kw at this temperature?
6.1 × 10⁻¹³
At 15°C, the value of Kw is 4.5 × 10⁻¹⁵. What is the equilibrium concentration of OH⁻ at this temperature?
6.7 × 10⁻⁸ M
A 0.100 M solution of weak acid HA ionizes 2.00%. What will be its percent ionization if 10.0 mL of this solution is diluted with 110.0 mL of water?
6.76 %
What is the entropy when 1.08 moles of CCl₂F₂ vaporize at 25°C? [∆H(vap) = 17.2 kJ/mol at 25°C]
62.3 J/K
How much heat would need to be removed to cool 150.3 g of water from 25.6°C to -10.7°C?
69.6 kJ
Complete and balance the following redox reaction in basic solution
7 IO₃⁻(aq) + 4 Re(s) + 4 OH⁻(aq) → 4 ReO₄⁻(aq) + 7 IO⁻(aq) + 2 H₂O(l)
How many grams of NH₄Br must be dissolved in 1.00 L of water to produce a solution with pH = 5.16? The Kb of NH₃ is equal to 1.8 × 10⁻⁵.
8.3g
A concentration cell was made using a silver metal electrode in a saturated solution of AgI for the anode half cell, and a silver metal electrode in 1.00 M Ag⁺ for the cathode half cell. E for the cell was 0.475 V at 25°C. Calculate the Ksp for AgI
8.58 × 10⁻¹⁷
Consider the equilibrium system described by the chemical reaction below. At equilibrium, a sample of gas from the system is collected into 5.00 L flask. The flask is found to contain 8.62 g of CO, 2.60 g of H₂, 43.0 g of CH₄, and 48.4 g of H₂O at 320 °C. What are the values of Kc and Kp for this reaction?
8.69
Determine the pH of a 0.65 M solution of HCO₂K. The value of Ka for HCOOH is 1.8 × 10⁻⁴.
8.78
Determine the pH at the equivalence point in the titration of 50.0 mL of 0.300 M CH₃COOH with 0.300 M NaOH. The value of Ka for CH₃COOH is 1.8 × 10⁻⁵.
8.96
For the equilibrium A + B ⇌ C + D K was measured to be 1.51 × 10⁻² at 25°C and 135.0 at 600°C. What is ∆S° for the reaction? Assume ∆H° and ∆S° are temperature independent.
80.0 J/mol･K
When heated, CaCO₃ decomposes into CaO and CO₂ according to the following equation: CaCO₃(s) ⇌ CaO(s) + CO₂(g) This is an endothermic reaction in which ∆H° = 178.3 kJ/mol and in which entropy increases, ∆S° = 159 J/mol･K. To what temperature (in °C) must CaCO₃ be heated in a closed container in order to produce CO₂ at an equilibrium pressure of 0.650 atm? (Assume ∆H° and ∆S° do not change appreciably with temperature).
824 °C
When heated, CaCO₃ decomposes into CaO and CO₂ according to the following equation: CaCO₃(s) ⇌ CaO(s) + CO₂(g) This is an endothermic reaction in which ∆H° = 178.3 kJ/mol and in which entropy increases, ∆S° = 159 J/mol･K. To what temperature (in °C) must CaCO₃ be heated in a closed container in order to produce CO₂ at an equilibrium pressure of 0.900 atm? (Assume ∆H° and ∆S° do not change appreciably with temperature).
842 °C
A 0.1000 M solution of a weak acid, HA, is 3.0% dissociated. Determine the value of Ka for the weak acid.
9.3 × 10^-5
Determine the pH at the point in the titration of 40.0 mL of 0.200 M H₂NNH₂ with 0.100 M HNO₃ after 10.0 mL of the strong acid has been added. The value of Kb for H₂NNH₂ is 3.0 × 10⁻⁶.
9.32
Determine the pH of the resulting solution when the following two solutions are mixed: 40.0 mL of 0.500 M NH₃ and 25.0 mL of 0.300 M HCl. The value of Kb for NH₃ is 1.8 × 10⁻⁵.
9.48
How many grams of copper metal can be deposited from Cu²⁺(aq) when a current of 2.50 A is run for 3.25 h? (F = 96,500 C/mol)
9.63 g
Calculate ∆E when 2 moles of a liquid is vaporized at its boiling point (85°C) and 1.00 atm pressure. ∆Hvap for the liquid is 48.1 kJ/mol at 85°C. Recall ∆H = ∆E + P∆V and watch your units!
90.2 kJ
Which of the following statements is true?
A buffer forms when a conjugate weak acid/weak base pair are mixed together.
A state function is best described as
A function that depends only on the starting point and ending point of a process
In which one of the following processes is the system endothermic?
A process taking place in solution where the solution temperature decreases
Which one of the following statements is true?
A solution with a pH of 8.20 is only slightly basic
The dissolution of ammonium nitrate in aqueous solution is an endothermic process and with large enough amounts of solid ammonium nitrate reaches an equilibrium. NH₄NO₃ (s) ⇌ NH₄⁺ (aq) + NO₃⁻ (aq) How will each of the following changes to a system at equilibrium affect the solubility of ammonium nitrate?
Adding solid ammonium nitrate : The solubility does not change Increasing the pH of the solution : The solubility increases Adding nitrate ion : The solubility decreases Increasing the temperature : The solubility increases
Which of the following pairs of substances has the species with the higher entropy listed first?
Ar(g), Hg(l)
Consider the following energy diagram and determine which of the following statements is true.
At equilibrium, we expect [Reactants] < [Products]
Which of the following titration curves represents the results for a weak acid titrated with a strong base?
B
A 0.012-mol sample of Na₂SO₄ is added to 400 mL of each of two solutions. One solution contains 1.5 × 10⁻³ M BaCl₂; the other contains 1.5 × 10⁻³ M CaCl₂. Ksp for BaSO₄ = 1.5 × 10⁻⁹ and Ksp for CaSO₄ = 6.1 × 10⁻⁵. Which of the following statements is true?
BaSO₄ would precipitate but CaSO₄ would not.
A heated gas expands and presses on a piston. In this example, work is done:
By the gas on the piston
The approximate pKa values of several acids are given in the table below. In which of the following acid-base reactions will the reverse reaction be favored?
CH₃OH(aq) + NH₃(aq) ⇌ CH₃O⁻(aq) + NH₄⁺(aq)
The following three equations represent equilibria that lie far to the right. HNO₃(aq) + CN⁻(aq) <--> HCN(aq) + NO₃⁻(aq) HCN(aq) + OH⁻(aq) <--> H₂O(l) + CN⁻(aq) H₂O(l) + CH₃O⁻(aq) <--> CH₃OH(aq) + OH⁻(aq) Identify the strongest base.
CH₃O⁻
For which one of the following is the enthalpy of the reaction the same as the enthalpy of formation?
Ca (s) + ½ O₂ (g) → CaO (s) (another answer could be 2 Fe (s) + ³/₂ O₂ (g) → Fe₂O₃ (s))
Determine if the following salt is neutral, acidic or basic. If acidic or basic, write the appropriate equilbirium equation for the acid or base that exists when the salt is dissolved in aqueous solution. If neutral, write only NR. RbClO
ClO⁻(aq) + H₂O(l) ⇌ HClO(aq) + OH⁻(aq)
A voltaic cell is prepared where copper metal is oxidized to copper(II) and silver(I) is reduced to silver metal. Which of the following reactions occurs at the anode?
Cu(s) → Cu²⁺(aq) + 2e⁻
Which of the following solutions would be expected to have a pH greater than 7.00?
C₆H₅COONa
A 0.200 M solution of XBr has a pH of 2.67. Which of the following could be XBr?
C₆H₅NH₃Br
Consider the table of weak bases below. Which of the following 1.0 M solutions would have the lowest pH?
C₆H₅NH₃⁺
The two salts AgX and AgY have very similar solubilities in water. The salt AgX is much more soluble in acid than is AgY. What can be said about the relative strengths of the acids HX and HY?
HY is stronger than HX.
Which of the following species is present in the greatest concentration in a 0.100 M H₂SO₄ solution in H₂O?
H₃O⁺
Determine if the following salt is neutral, acidic or basic. If acidic or basic, write the appropriate equilibrium equation for the acid or base that exists when the salt is dissolved in aqueous solution. If neutral, write only NR.
IO⁻(aq) + H₂O(l) ⇌ HIO(aq) + OH⁻(aq)
Write the basic equilibrium equation for IO⁻
IO⁻(aq) + H₂O(l) ⇌ OH⁻(aq) + HIO(aq)
Consider the equilibrium system described by the chemical reaction below. A mixture of gas containing only N₂ and H₂ is reacted in a vessel at high temperature. At equilibrium, the 5.0 M H₂, 8.0 M N₂, and 4.0 M NH₃ are present. Determine the initial concentrations of H₂ and N₂ that were present in the vessel.
Ice table
Which of the following accurately describes the species in solution at point C on the titration curve for the titration of NH₃ with HCl?
NH₄⁺
Which of the following 1.0 M solutions would have the highest pH?
NaIO
Which one of the following titrations is expected to have a pH < 7 at the equivalence point?
NaNO₂ titrated with HBr
Which of the following shows what will happen when Na₂O is dissolved in water?
Na₂O (s) + H₂O (l) ⇌ 2 NaOH (aq)
A solution is made by mixing 3.5 × 10⁻⁴ M Sr²⁺ and 0.0010 M F⁻. Will a precipitate form? (Ksp of SrF₂ is 4.3 × 10⁻⁹)
No, because Q < Ksp
A decrease of pH by 3 implies
The OH⁻ concentration decreases by a factor of 1000
Which of the following statements is incorrect about a concentration cell consisting of the same metal and metal ion solutions at different concentrations?
The cell containing the more concentrated solution is the anode
Given the standard electrode reduction potentials, which reaction will occur spontaneously?
V + 2Cu²⁺ → V²⁺ + 2Cu⁺
(On phone) A 4.99 g sample of an unknown salt (MM = 116.82 g/mol) is dissolved in 150.00 g water in a coffee cup calorimeter. Before placing the sample in the water, the temperature of the salt and water is 23.72 °C. After the salt has completely dissolved, the temperature of the solution is 28.54 °C.
What is the total mass inside the calorimeter? : 154.99 g What is the change in temperature inside the calorimeter? : 4.82 °C Based on a temperature change of 4.82 °C, was the dissolution process endothermic or exothermic? : Exothermic because the water temperature increased How much heat was gained if 154.99 g of solution increased in temperature by 4.82 °C? Assume the specific heat of the solution is the same as water, 4.184 J/g･°C. : 3126 J If 3126 J of heat was gained by the solution, what is the heat for the dissolution reaction? : -3126 J how many moles of the unknown salt were used in the reaction? : 0.0427 mol If -3126 J of heat was lost during the dissolution reaction of 0.0427 moles of the unknown salt, what is the enthalpy change (in kJ/mol) for the dissolution reaction? : -73.2 kJ/mol
The weak acid HY is much stronger than weak acid HX. Which one of the following statements is true?
Y⁻ is a weaker base than X⁻
A solution is made by dissolving 28.7 g of CH₃NH₃NO₃ in 500.0 mL of water.
answers on phone
Consider a buffer made by adding 38.3 g of (CH₃)₂NH₂I to 250.0 mL of 1.42 M (CH₃)₂NH (Kb = 5.4 x 10⁻⁴)
answers on phone
Consider the titration of a 60.0 mL of 0.279 M weak acid HA (Ka = 4.2 x 10⁻⁶) with 0.400 M KOH.
answers on phone
Based on the law of conservation of energy, which of the following describes the transformation of energy in a diesel engine?
chemical energy to kinetic energy and heat
An equilibrium is established for the exothermic reaction Br₂(g) + 5 F₂(g) ⇌ 2 BrF₅(g). How would each of the following changes affect the partial pressure of fluorine at equilibrium?
on my phone
An equilibrium is established for the reaction 2 CO(g) + MoO₂(s) ⇌ 2 CO₂(g) + Mo(s).
on my phone
Consider a buffer made by adding 132.8 g of NaC₇H₅O₂ to 300.0 mL of 1.15 M HC₇H₅O₂ (Ka = 6.3 x 10⁻⁵)
on my phone
Consider the following equation in aqueous solution: PO₃³⁻(aq) + SO₂(g) → PO₄³⁻(aq) + S²⁻(aq)
on my phone
Consider the titration of a 40.0 mL of 0.219 M weak acid HA (Ka = 2.7 x 10⁻⁸) with 0.100 M LiOH.
on my phone
The value of Kp for the reaction 2 A(g) + B(g) + 3 C(g) → 2 D(g) + E(g) is 15570 at a particular temperature.
on my phone
Which of the following is NOT a form of potential energy?
thermal
An exothermic reaction causes the surroundings to
warm up
Which of the following is true for a particular reaction if ∆G° is -40.0 kJ/mol at 290 K and -20.0 kJ/mol at 390 K?
∆H < 0, ∆S < 0
Which of the following sets of conditions is true for an exothermic reaction that is spontaneous at all temperatures?
∆H < 0, ∆S > 0, ∆G < 0
Liquid water at 50°C is heated to steam at 150°C. Which piece of information is NOT required to determine the amount of heat that must be added for this change to occur?
∆H(fusion)
Which of the following is true for all reversible exothermic processes?
∆H(sys) < 0, ∆S(surr) > 0
Which of the following guarantees a reaction will be spontaneous?
∆S > ∆H/T
Which of the following is true for a system at equilibrium?
∆S°(sys) = -∆S°(surr)
What is Kb for the conjugate base of HNO₂ (Ka = 4.5 × 10⁻⁴)?
2.2 × 10⁻¹¹
At what temperature, in K, does X(l) → X(g) occur spontaneously?
349 K
Which of the following processes is both exothermic (at 1 atm) and spontaneous?
Condensation of water at 99.5°C
One hour of bicycle riding can require 500-900 kcal of energy, depending on the speed, the terrain, and the weight of the racer. An individual wants to lose weight by riding at a modest effort, consuming 505 kcal of energy per hour. How many hours must he ride to lose 21.0 pounds of body fat? One gram of body fat is equivalent to 7.70 kcal of energy. There are 454 g in 1 lb.
145 hr
The concentration of Al³⁺ in a saturated solution of Al(OH)₃ at 25°C is 5.2 × 10⁻⁹ M. Calculate the Ksp for Al(OH)₃
2.0 × 10⁻³²
At what pH will Fe(OH)₃(s) begin to precipitate from 2.35 × 10⁻³ M FeCl₃? Ksp for Fe(OH)₃ is 2.79 × 10⁻³⁹
2.02
What is the equilibrium constant for the solubility of FeCO₃ (Ksp = 2.1 × 10⁻¹¹) in NaCN? (Kf of Fe(CN)₆⁴⁻ is 1.0 × 10³⁵)
2.1 × 10²⁴
What is the solubility of PbF₂ in a solution that contains 0.0400 M F⁻ ions? (Ksp of PbF₂ is 3.60 × 10⁻⁸)
2.25 × 10⁻⁵ M
What is the solubility of AgCN(s) (Ksp = 2.2 × 10⁻¹⁶) in a solution containing 0.25 M H⁺? (Ka for HCN is 6.2 × 10⁻¹⁰.)
2.98 × 10⁻⁴ M
The combustion of propane (C₃H₈) produces 2220 kJ of energy per mole of propane consumed. How many grams of propane will be required to heat 55.0 gal of bathtub water from 25.0°C to 35.0°C if the process is 80.0% efficient? (1 gal = 3.785 L, 1 cal = 4.184 J, the density of water is 1.00 g/mL, the specific heat of water is 1 cal/(g°C)
216 g
Compared to HCO₃⁻, H₂CO₃ has a
Higher Ka; Lower pKa
Which of the following processes has a negative ∆S?
H₂O(g) → H₂O(s)
How does the third law of thermodynamics allow absolute entropies of substances to be determined?
It defines a reference point by which entropy changes can be measured and assigned as an absolute entropy for a substance.
The reaction Q(g) + R(g) → Z(l) is shown to be exothermic. Which of the following is true concerning the reaction?
It is spontaneous only at low T
When boiling water on a natural gas (methane) burner, which of the following processes converts chemical potential energy to heat?
The combustion reaction between methane and oxygen
Which of the following is true for a reaction when ∆G is a positive value?
The reverse of the reaction is spontaneous
When water cools from 10 °C and then freezes to become ice, which of the following best describes describe the heat flow between the system and surroundings?
The water is releasing energy as its temperature decreases, and that energy is absorbed by the surroundings
Salt A has a greater solubility in water than Salt B. What can be said about their Ksp values?
Their Ksp values cannot be compared because we do not know the number of ions each produces
