Question 1

When 50.0 mL of 0.400 M Ca(NO₃)₂ is added to 50.0 mL of 0.800 M NaF,CaF₂ precipitates,as shown in the net ionic equation below.The initial temperature of both solutions is 23.0°C.Assuming that the reaction goes to completion,and that the resulting solution has a mass of 100.00 g and a specific heat of 4.18 J/(g ∙ °C),calculate the final temperature of the solution. Ca²⁺(aq)+ 2 F^-(aq)→ CaF₂(s)ΔH° = -11.5 kJ

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Question 2

At 1 atm pressure,the heat of sublimation of gallium is 277 kJ/mol and the heat of vaporization is 271 kJ/mol.To the correct number of significant figures,how much heat is required to melt 4.50 mol of gallium at 1 atm pressure?

Accepted Answer

The heat of fusion is the difference between the heat of sublimation and the heat of vaporization. Therefore, the heat required to melt 4.50 mol of gallium is:$$\Delta H_{fusion} = \Delta H_{sublimation} - \Delta H_{vaporization}$$$$\Delta H_{fusion} = 277 kJ/mol - 271 kJ/mol = 6 kJ/mol$$Therefore, the heat required to melt 4.50 mol of gallium is:$$\Delta H_{fusion} = 6 kJ/mol * 4.50 mol = 27 kJ$$

Question 3

How much heat is absorbed/released when 35.00 g of NH₃(g)reacts in the presence of excess O₂(g)to produce NO(g)and H₂O(l)according to the following chemical equation? 4 NH₃(g)+ 5 O₂(g)→ 4 NO(g)+ 6 H₂O(l)ΔH° = 1168 kJ

Accepted Answer

We can use the given equation and the given value of ΔH° to calculate the amount of heat absorbed or released in the reaction. First, we need to convert the given mass of NH₃(g) to moles: 35.00 g NH₃(g) x (1 mol NH₃(g) / 168.15 g NH₃(g)) = 0.2082 mol NH₃(g) According to the balanced equation, 4 moles of NH₃(g) produce 1168 kJ of heat. Therefore, 0.2082 moles of NH₃(g) will produce: 1168 kJ x (0.2082 mol / 4 mol) = 60.01 kJ of heat The reaction produces NO(g) and H₂O(l) from NH₃(g) and O₂(g), so it is an endothermic reaction that absorbs heat. Therefore, the correct choice is A) 600.1 kJ of heat are absorbed.

Question 4

Sodium metal reacts with water to produce hydrogen gas and sodium hydroxide according to the chemical equation shown below.When 0.025 mol of Na is added to 100.00 g of water,the temperature of the resulting solution rises from 25.00°C to 35.75°C.If the specific heat of the solution is 4.18 J/(g ∙ °C),calculate ΔH for the reaction,as written. 2 Na(s)+ 2 H₂O(l)→ 2 NaOH(aq)+ H₂(g)ΔH= ?

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Question 5

In the presence of excess oxygen,methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH₄(g)+ 2O₂(g)→ CO₂(g)+ 2H₂O(l)ΔH = -890.0 kJ Calculate the value of q (kJ)in this exothermic reaction when 1.70 g of methane is combusted at constant pressure.

Accepted Answer

First, we need to calculate the amount of methane that is being combusted. We are given the mass of methane, which we can convert to moles using its molar mass:

1.70 g CH4 × (1 mol CH4/16.04 g CH4) = 0.106 mol CH4

Next, we can use the balanced chemical equation to relate the moles of methane consumed to the heat released:

0.106 mol CH4 × (-890.0 kJ/1 mol CH4) = -94.6 kJ

Note that the negative sign indicates that heat is released from the system, which makes sense since this is an exothermic reaction. Therefore, the correct answer is A.

Question 6

According to the following thermochemical equation,what mass of H₂O (in g)must form in order to produce 488 kJ of energy? SiO₂(s)+ 4 HF(g)→ SiF₄(g)+ 2 H₂O(l)ΔH°_rxn= -184 kJ

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Question 7

The combustion of titanium with oxygen produces titanium dioxide: Ti(s)+ O₂(g)→ TiO₂(s) When 2.060 g of titanium is combusted in a bomb calorimeter,the temperature of the calorimeter increases from 25.00°C to 91.60°C.In a separate experiment,the heat capacity of the calorimeter is measured to be 9.84 kJ/K.The heat of reaction for the combustion of a mole of Ti in this calorimeter is ________ kJ/mol.

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Question 8

When 1.50 g of Ba(s)is added to 100.00 g of water in a container open to the atmosphere,the reaction shown below occurs and the temperature of the resulting solution rises from 22.00°C to 33.10°C.If the specific heat of the solution is 4.18 J/(g ∙ °C),calculate ΔH for the reaction,as written. Ba(s)+ 2 H₂O(l)→ Ba(OH)₂(aq)+ H₂(g)ΔH = ?

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Question 9

Hydrogen peroxide decomposes to water and oxygen at constant pressure by the following reaction: 2H₂O₂(l)→ 2H₂O(l)+ O₂(g)ΔH = -196 kJ Calculate the value of q (kJ)in this exothermic reaction when 4.00 g of hydrogen peroxide decomposes at constant pressure?

Accepted Answer

The molar mass of H₂O₂ is 34.01 g/mol. 4.00 g of H₂O₂ is 4.00 g / 34.01 g/mol = 0.1176 mol. The reaction releases 196 kJ for 2 moles of H₂O₂, so for 0.1176 mol, q = (196 kJ / 2 mol) * 0.1176 mol = -11.5 kJ.

Question 10

Which of the following substances (with specific heat capacity provided)would show the greatest temperature change upon absorbing 100.0 J of heat?

Accepted Answer

The substance with the lowest specific heat capacity will show the greatest temperature change for the same amount of heat absorbed. Lead (Pb) has the lowest specific heat capacity (0.128 J/g°C) among the given options.

Question 11

For a particular process that is carried out at constant pressure,q = 145 kJ and w = -35 kJ.Therefore,

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Question 12

When 10.00 moles of H₂(g)reacts with 5.000 mol of O₂(g)to form 10.00 mol of H₂O(l)at 25°C and a constant pressure of 1.00 atm.If 683.0 kJ of heat are released during this reaction,and PΔV is equal to -37.00 kJ,then

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Question 13

What is the enthalpy change (in kJ)of a chemical reaction that raises the temperature of 250.0 ml of solution having a density of 1.25 g/ml by 7.80°C? (The specific heat of the solution is 3.74 j/g ∙ K. )

Accepted Answer

Firstly, we need to calculate the mass of the solution:

mass = volume x density = 250.0 ml x 1.25 g/ml = 312.5 g

Now, we can use the formula:

ΔH = m x c x ΔT

where: 
m = mass of the solution (in grams) = 312.5 g 
c = specific heat of the solution = 3.74 J/g*K 
ΔT = change in temperature = 7.80°C

ΔH = 312.5 g x 3.74 J/g*K x 7.80°C = 9152.25 J = 9.15 kJ

Therefore, the enthalpy change of the chemical reaction is -9.15 kJ (since the temperature increase is due to an exothermic reaction). Rounded to two decimal places, the answer is D) -9.12 kJ.

Question 14

At constant pressure,the combustion of 15.0 g of C₂H₆(g)releases 777 kJ of heat.Calculate ΔH for the reaction given below. 2 C₂H₆(g)+ 7 O₂(g)→ 4 CO₂(g)+ 6 H₂O(l)

Accepted Answer

In order to calculate ΔH for the given reaction, we need to use the equation: ΔH = Σ nΔH(products) - Σ nΔH(reactants) where ΣnΔH is the sum of the enthalpies of formation of the products or reactants, multiplied by their stoichiometric coefficients. We can use the following values for the enthalpies of formation (in kJ/mol): CO₂(g) = -107.1 H₂O(l) = -285.8 C₂H₆(g) = 109.5 O₂(g) = 0 Using these values and the stoichiometric coefficients from the balanced equation, we can calculate: ΔH = (4 x -107.1 kJ/mol + 6 x -285.8 kJ/mol) - (2 x 109.5 kJ/mol + 7 x 0 kJ/mol) ΔH = -3120 kJ/mol Therefore, the answer is D, -3120 kJ.

Question 15

How much heat is absorbed when 45.00 g of C(s)reacts in the presence of excess SO₂(g)to produce CS₂(l)and CO(g)according to the following chemical equation? 5 C(s)+ 2 SO₂(g)→ CS₂(l)+ 4 CO(g)ΔH° = 239.9 kJ

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Question 16

When 5.00 mol of benzene is vaporized at a constant pressure of 1.00 atm and at its normal boiling point of 80.1°C,169.5 kJ are absorbed and PΔV for the vaporization process is equal to 14.5 kJ then

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Question 17

When 0.455 g of anthracene,C₁₄H₁₀,is combusted in a bomb calorimeter that has a water jacket containing 500.0 g of water,the temperature of the water increases by 8.63°C.Assuming that the specific heat of water is 4.18 J/(g ∙ °C),and that the heat absorption by the calorimeter is negligible,estimate the enthalpy of combustion per mole of anthracene.

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Question 18

A 50.0-g sample of liquid water at 25.0°C is mixed with 29.0 g of water at 45.0°C.The final temperature of the water is ________°C.

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Question 19

When 1.50 mol of CH₄(g)reacts with excess Cl₂(g)at constant pressure according to the chemical equation shown below,1062 kJ of heat are released.Calculate the value of ΔH for this reaction,as written. 2 CH₄(g)+ 3 Cl₂(g)→ 2 CHCl₃(l)+ 3 H₂(g)ΔH = ?

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Question 20

Identify an energy source that is not renewable.

Accepted Answer

Nuclear energy is not renewable because it relies on the use of uranium, a finite resource, and the process of nuclear fission. Once the uranium has been used up, it cannot be replaced, making nuclear energy non-renewable.

Quiz 6: Thermochemistry

At 1 atm pressure,the heat of sublimation of gallium is 277 kJ/mol and the heat of vaporization is 271 kJ/mol.To the correct number of significant figures,how much heat is required to melt 4.50 mol of gallium at 1 atm pressure?

How much heat is absorbed/released when 35.00 g of NH₃(g) reacts in the presence of excess O₂(g) to produce NO(g) and H₂O(l) according to the following chemical equation? 4 NH₃(g) + 5 O₂(g) → 4 NO(g) + 6 H₂O(l) ΔH° = 1168 kJ

In the presence of excess oxygen,methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = -890.0 kJ Calculate the value of q (kJ) in this exothermic reaction when 1.70 g of methane is combusted at constant pressure.

According to the following thermochemical equation,what mass of H₂O (in g) must form in order to produce 488 kJ of energy? SiO₂(s) + 4 HF(g) → SiF₄(g) + 2 H₂O(l) ΔH°_rxn= -184 kJ

Hydrogen peroxide decomposes to water and oxygen at constant pressure by the following reaction: 2H₂O₂(l) → 2H₂O(l) + O₂(g) ΔH = -196 kJ Calculate the value of q (kJ) in this exothermic reaction when 4.00 g of hydrogen peroxide decomposes at constant pressure?

Which of the following substances (with specific heat capacity provided) would show the greatest temperature change upon absorbing 100.0 J of heat?

For a particular process that is carried out at constant pressure,q = 145 kJ and w = -35 kJ.Therefore,

When 10.00 moles of H₂(g) reacts with 5.000 mol of O₂(g) to form 10.00 mol of H₂O(l) at 25°C and a constant pressure of 1.00 atm.If 683.0 kJ of heat are released during this reaction,and PΔV is equal to -37.00 kJ,then

What is the enthalpy change (in kJ) of a chemical reaction that raises the temperature of 250.0 ml of solution having a density of 1.25 g/ml by 7.80°C? (The specific heat of the solution is 3.74 j/g ∙ K. )

At constant pressure,the combustion of 15.0 g of C₂H₆(g) releases 777 kJ of heat.Calculate ΔH for the reaction given below. 2 C₂H₆(g) + 7 O₂(g) → 4 CO₂(g) + 6 H₂O(l)

How much heat is absorbed when 45.00 g of C(s) reacts in the presence of excess SO₂(g) to produce CS₂(l) and CO(g) according to the following chemical equation? 5 C(s) + 2 SO₂(g) → CS₂(l) + 4 CO(g) ΔH° = 239.9 kJ

When 5.00 mol of benzene is vaporized at a constant pressure of 1.00 atm and at its normal boiling point of 80.1°C,169.5 kJ are absorbed and PΔV for the vaporization process is equal to 14.5 kJ then

A 50.0-g sample of liquid water at 25.0°C is mixed with 29.0 g of water at 45.0°C.The final temperature of the water is ________°C.

When 1.50 mol of CH₄(g) reacts with excess Cl₂(g) at constant pressure according to the chemical equation shown below,1062 kJ of heat are released.Calculate the value of ΔH for this reaction,as written. 2 CH₄(g) + 3 Cl₂(g) → 2 CHCl₃(l) + 3 H₂(g) ΔH = ?

Identify an energy source that is not renewable.

Quiz 6: Thermochemistry

At 1 atm pressure,the heat of sublimation of gallium is 277 kJ/mol and the heat of vaporization is 271 kJ/mol.To the correct number of significant figures,how much heat is required to melt 4.50 mol of gallium at 1 atm pressure?

How much heat is absorbed/released when 35.00 g of NH3(g) reacts in the presence of excess O2(g) to produce NO(g) and H2O(l) according to the following chemical equation? 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(l) ΔH° = 1168 kJ

In the presence of excess oxygen,methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = -890.0 kJ Calculate the value of q (kJ) in this exothermic reaction when 1.70 g of methane is combusted at constant pressure.

According to the following thermochemical equation,what mass of H2O (in g) must form in order to produce 488 kJ of energy? SiO2(s) + 4 HF(g) → SiF4(g) + 2 H2O(l) ΔH°rxn = -184 kJ

Hydrogen peroxide decomposes to water and oxygen at constant pressure by the following reaction: 2H2O2(l) → 2H2O(l) + O2(g) ΔH = -196 kJ Calculate the value of q (kJ) in this exothermic reaction when 4.00 g of hydrogen peroxide decomposes at constant pressure?

Which of the following substances (with specific heat capacity provided) would show the greatest temperature change upon absorbing 100.0 J of heat?

For a particular process that is carried out at constant pressure,q = 145 kJ and w = -35 kJ.Therefore,

When 10.00 moles of H2(g) reacts with 5.000 mol of O2(g) to form 10.00 mol of H2O(l) at 25°C and a constant pressure of 1.00 atm.If 683.0 kJ of heat are released during this reaction,and PΔV is equal to -37.00 kJ,then

What is the enthalpy change (in kJ) of a chemical reaction that raises the temperature of 250.0 ml of solution having a density of 1.25 g/ml by 7.80°C? (The specific heat of the solution is 3.74 j/g ∙ K. )

At constant pressure,the combustion of 15.0 g of C2H6(g) releases 777 kJ of heat.Calculate ΔH for the reaction given below. 2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l)

How much heat is absorbed when 45.00 g of C(s) reacts in the presence of excess SO2(g) to produce CS2(l) and CO(g) according to the following chemical equation? 5 C(s) + 2 SO2(g) → CS2(l) + 4 CO(g) ΔH° = 239.9 kJ

When 5.00 mol of benzene is vaporized at a constant pressure of 1.00 atm and at its normal boiling point of 80.1°C,169.5 kJ are absorbed and PΔV for the vaporization process is equal to 14.5 kJ then

A 50.0-g sample of liquid water at 25.0°C is mixed with 29.0 g of water at 45.0°C.The final temperature of the water is ________°C.

When 1.50 mol of CH4(g) reacts with excess Cl2(g) at constant pressure according to the chemical equation shown below,1062 kJ of heat are released.Calculate the value of ΔH for this reaction,as written. 2 CH4(g) + 3 Cl2(g) → 2 CHCl3(l) + 3 H2(g) ΔH = ?

Identify an energy source that is not renewable.

How much heat is absorbed/released when 35.00 g of NH₃(g) reacts in the presence of excess O₂(g) to produce NO(g) and H₂O(l) according to the following chemical equation? 4 NH₃(g) + 5 O₂(g) → 4 NO(g) + 6 H₂O(l) ΔH° = 1168 kJ

In the presence of excess oxygen,methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = -890.0 kJ Calculate the value of q (kJ) in this exothermic reaction when 1.70 g of methane is combusted at constant pressure.

According to the following thermochemical equation,what mass of H₂O (in g) must form in order to produce 488 kJ of energy? SiO₂(s) + 4 HF(g) → SiF₄(g) + 2 H₂O(l) ΔH°_rxn= -184 kJ

Hydrogen peroxide decomposes to water and oxygen at constant pressure by the following reaction: 2H₂O₂(l) → 2H₂O(l) + O₂(g) ΔH = -196 kJ Calculate the value of q (kJ) in this exothermic reaction when 4.00 g of hydrogen peroxide decomposes at constant pressure?

When 10.00 moles of H₂(g) reacts with 5.000 mol of O₂(g) to form 10.00 mol of H₂O(l) at 25°C and a constant pressure of 1.00 atm.If 683.0 kJ of heat are released during this reaction,and PΔV is equal to -37.00 kJ,then

At constant pressure,the combustion of 15.0 g of C₂H₆(g) releases 777 kJ of heat.Calculate ΔH for the reaction given below. 2 C₂H₆(g) + 7 O₂(g) → 4 CO₂(g) + 6 H₂O(l)

How much heat is absorbed when 45.00 g of C(s) reacts in the presence of excess SO₂(g) to produce CS₂(l) and CO(g) according to the following chemical equation? 5 C(s) + 2 SO₂(g) → CS₂(l) + 4 CO(g) ΔH° = 239.9 kJ

When 1.50 mol of CH₄(g) reacts with excess Cl₂(g) at constant pressure according to the chemical equation shown below,1062 kJ of heat are released.Calculate the value of ΔH for this reaction,as written. 2 CH₄(g) + 3 Cl₂(g) → 2 CHCl₃(l) + 3 H₂(g) ΔH = ?