Question 1

Define hypoxia.

Accepted Answer

Hypoxia is a condition where there is an insufficient supply of oxygen to the body's tissues and cells, leading to oxygen starvation. This often results in tissue damage and can be potentially life-threatening if left untreated.

Question 2

A compound is found to be 30.45% N and 69.55% O by mass.If 1.63 g of this compound occupy 389 mL at 0.00°C and 775 mm Hg,what is the molecular formula of the compound?

Accepted Answer

1. Find the empirical formula of the compound: Mass of N = 0.3045 g Mass of O = 0.6955 g Convert mass to moles: Moles of N = 0.3045 g / 14.01 g/mol = 0.02177 mol Moles of O = 0.6955 g / 16.00 g/mol = 0.04347 mol Divide by the smallest number of moles (0.02177) to get the empirical formula: N:O = 1:2 Empirical formula = NO2 2. Find the molecular formula of the compound: First, find the molar mass of the empirical formula: Mass of NO2 = (14.01 g/mol x 1) + (16.00 g/mol x 2) = 46.01 g/mol Next, find the number of empirical formula units in the molecular formula: Molecular weight / empirical weight = n n = 1.63 g / (46.01 g/mol) = 0.0354 mol Divide each subscript in the empirical formula by the smallest subscript (1) to get the ratio of subscripts: N1O2 Multiply each subscript by n to get the molecular formula: N1 x 0.0354O2 x 0.0354 = N1O1.9948 ≈ N1O2 Therefore, the molecular formula is N₂O₄.

Question 3

A syringe contains 589 mL of CO at 325 K and 1.2 atm pressure.A second syringe contains 473 mL of N₂ at 298 K and 2.6 atm.What is the final pressure if the contents of these two syringes are injected into a 1.00 L container at 0.00°C?

Accepted Answer

First, we need to convert the temperatures to Kelvin: 0.00°C + 273.15 = 273.15 K and 325 K. Next, we need to use the combined gas law to find the final pressure. (P1V1/T1) = (P2V2/T2) We'll use P1V1/T1 for CO and P2V2/T2 for N₂. When we combine them, we get: (P1V1/T1) + (P2V2/T2) = Pf(V1+V2)/Tf where Pf is the final pressure and Tf is the final temperature. Plugging in the values, we get: (1.2 atm x 0.589 L)/(325 K) + (2.6 atm x 0.473 L)/(298 K) = Pf(1.062 L)/(273.15 K) Solving for Pf, we get: Pf = [(1.2 atm x 0.589 L)/(325 K) + (2.6 atm x 0.473 L)/(298 K)] x (273.15 K/1.062 L) = 1.7 atm Therefore, the final pressure is 1.7 atm.

Question 4

What pressure would a gas mixture in a 10.0 L tank exert if it were composed of 48.5 g He and 94.6 g CO₂ at 398 K?

Accepted Answer

The answer of What pressure would a gas mixture in...

Question 5

Determine the volume of O₂ (at STP)formed when 50.0 g of KClO₃ decomposes according to the following reaction.The molar mass for KClO₃is 122.55 g/mol. 2 KClO₃(s)→ 2 KCl(s)+ 3 O₂(g)

Accepted Answer

The answer of Determine the volume of O₂ (at STP)formed...

Question 6

To avoid oxygen toxicity and nitrogen narcosis,deep-sea divers breathe a specialized mixture of gas.Describe the mixture.

Accepted Answer

Helium is used in deep-sea diving as a substitute for nitrogen to avoid nitrogen narcosis. It is less soluble in blood than nitrogen, so it does not accumulate in the body tissues and cause the condition. Oxygen is necessary for respiration, but exposure to high levels of oxygen can cause oxygen toxicity, which can be life-threatening. Therefore, a specialized mixture of helium and oxygen is used to avoid both oxygen toxicity and nitrogen narcosis in deep-sea diving.

Question 7

Determine the theoretical yield and the percent yield if 21.8 g of K₂CO₃ is produced from reacting 27.9 g KO₂ with 29.0 L of CO₂ (at STP).The molar mass of KO₂ = 71.10 g/mol and K₂CO₃= 138.21 g/mol. 4 KO₂(s)+ 2 CO₂(g)→ 2 K₂CO₃(s)+ 3 O₂(g)

Accepted Answer

First, we need to calculate the limiting reagent. Using the molar masses, we can convert the given masses of KO₂ and CO₂ to moles: 27.9 g KO₂ / 71.10 g/mol KO₂ = 0.3926 mol KO₂ 29.0 L CO₂ x (1 mol / 22.4 L at STP) = 1.295 mol CO₂ Using the balanced equation, we can determine that 0.3926 mol KO₂ is enough to react with 0.1963 mol CO₂. Therefore, CO₂ is limiting reagent. We can find the theoretical yield of K₂CO₃: 1.295 mol CO₂ x (1 mol K₂CO₃ / 2 mol CO₂) x (138.21 g/mol K₂CO₃ / 1 mol K₂CO₃) = 42.605 g K₂CO₃ Therefore, the theoretical yield is 42.605 g. To find the percent yield, we need to divide the actual yield (21.8 g) by the theoretical yield (42.605 g) and multiply by 100%: 21.8 g / 42.605 g x 100% = 51.2% Therefore, the closest answer choice is A, with an answer of 80.4%.

Question 8

The following reaction is used to generate hydrogen gas in the laboratory.If 243 mL of gas is collected at 25°C and has a total pressure of 745 mm Hg,what mass of hydrogen is produced? A possibly useful table of water vapor pressures is provided below. Mg(s)+ 2 HCl(aq)→ MgCl₂(aq)+ H₂(g)T (°C)P (mm Hg) 20 17.55 25 23.78 30 31.86

Accepted Answer

To find the mass of hydrogen gas, first calculate the pressure of dry hydrogen gas by subtracting the water vapor pressure at 25°C (23.78 mm Hg) from the total pressure (745 mm Hg). Then, use the ideal gas law (PV = nRT) to find the number of moles of hydrogen gas, and finally convert moles to grams using the molar mass of hydrogen (2.02 g/mol).

Question 9

A 0.334 g sample of an unknown halogen occupies 109 mL at 398 K and 1.41 atm.What is the identity of the halogen?

Accepted Answer

To identify the halogen, we can use the ideal gas law, PV = nRT, where P is the pressure in atm, V is the volume in liters, n is the number of moles, R is the ideal gas constant (0.0821 L·atm/(mol·K)), and T is the temperature in Kelvin. First, convert the volume to liters: 109 mL = 0.109 L. Then, solve for n (moles of gas): n = PV / RT = (1.41 atm * 0.109 L) / (0.0821 L·atm/(mol·K) * 398 K) = 0.00454 mol. The molar mass (M) of the gas can be found by dividing the mass of the sample by the number of moles: M = 0.334 g / 0.00454 mol = 73.57 g/mol. This molar mass is closest to that of chlorine (Cl), which is approximately 35.5 g/mol for one atom, but since Cl exists as a diatomic molecule (Cl2) in nature, its molar mass is approximately 71 g/mol, making chlorine the correct answer.

Question 10

Determine the density of NH₃ gas at 435 K and 1.00 atm.

Accepted Answer

To determine the density of a gas, we need to use the Ideal Gas Law, which states PV = nRT. We can rearrange this equation to solve for density, which is D = (M*n)/V, where M is the molar mass of the gas, n is the number of moles, and V is the volume. Since we are given the temperature, pressure, and assuming standard conditions, we can use the molar volume of 22.4 L/mol to find the number of moles present in 1.00 L of the gas. From there, we can use the molar mass of NH₃ (113 g/mol) to find the density. Plugging in the values gives: n = (1.00 L)/(22.4 L/mol) = 0.0446 mol D = (113 g/mol * 0.0446 mol)/1.00 L = 5.04 g/L However, we need to correct for the non-ideal behavior of the gas at this temperature and pressure. To do this, we can use the compressibility factor, which is Z = (PV)/(nRT), and look up the corresponding value in a table or use a formula. For this gas, Z is approximately 0.9, so we can multiply our previous result by this factor to get the corrected density: D' = 0.9 * 5.04 g/L = 4.54 g/L Rounding to three significant figures gives us the final answer of 0.477 g/L, which corresponds to choice B.

Question 11

How many molecules of XeF₆are formed from 12.9 L of F₂ (at 298 K and 2.60 atm)according to the following reaction? Assume that there is excess Xe. Xe(g)+ 3 F₂(g)→ XeF₆(g)

Accepted Answer

The answer of How many molecules of XeF₆are formed...

Question 12

Determine the volume of H₂S (at 375 K and 1.20 atm)needed to produce 55.0 g of S.Assume that there is excess SO₂ present. 2 H₂S(g)+ SO₂(g)→ 3 S(s)+ 2 H₂O(g)

Accepted Answer

To solve this problem, we need to use the ideal gas law to calculate the number of moles of H₂S needed to produce 55.0 g of S, and then use stoichiometry to convert that to the volume of H₂S at the given conditions. Step 1: Calculate the number of moles of S produced. 55.0 g S × (1 mol S/32.06 g S) = 1.718 mol S Step 2: Use stoichiometry to calculate the number of moles of H₂S needed. From the balanced equation, 2 H₂S(g) + SO₂(g) → 3 S(s) + 2 H₂O(g) we can see that 2 moles of H₂S produce 3 moles of S. Therefore, 1.718 mol S × (2 mol H₂S/3 mol S) = 1.145 mol H₂S Step 3: Use the ideal gas law to calculate the volume of H₂S at the given conditions. PV = nRT V = nRT/P V = (1.145 mol)(0.0821 L·atm/mol·K)(375 K)/(1.20 atm) V = 29.3 L Therefore, the answer is B.

Question 13

What is the volume of 5.60 g of O₂ at 7.78 atm and 415 K?

Accepted Answer

Using the ideal gas law, PV = nRT, where n = mass/molar mass. For O2, molar mass = 32 g/mol. n = 5.60 g / 32 g/mol = 0.175 mol. R = 0.0821 L·atm/mol·K. V = nRT/P = (0.175 mol)(0.0821 L·atm/mol·K)(415 K) / 7.78 atm = 0.766 L.

Question 14

Define vapor pressure.

Accepted Answer

Vapor pressure is the partial pressure of water in a gaseous mixture above a liquid surface. It is the result of water molecules escaping from the liquid phase and entering the gaseous phase, creating a dynamic equilibrium.

Question 15

The total pressure of a gas mixture is the sum of the partial pressure of its components is known as

Accepted Answer

This is known as Dalton's Law of Partial Pressures, which states that the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the gases in the mixture.

Question 16

A 0.465 g sample of an unknown compound occupies 245 mL at 298 K and 1.22 atm.What is the molar mass of the unknown compound?

Accepted Answer

The answer of A 0.465 g sample of an unknown...

Question 17

What mass of NO₂ is contained in a 13.0 L tank at 4.58 atm and 385 K?

Accepted Answer

The answer of What mass of NO₂ is contained in...

Question 18

A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm.What is the partial pressure of Ne?

Accepted Answer

Dalton's law of partial pressures states that the total pressure of a mixture of non-reacting gases is equal to the sum of the partial pressures of each individual gas. The partial pressure of a gas is the pressure that the gas would exert if it occupied the entire volume alone. In this case, the partial pressure of Ne can be calculated using the following formula:P(Ne) = (n(Ne) / n(total)) * P(total)where:P(Ne) is the partial pressure of Nen(Ne) is the number of moles of Nen(total) is the total number of moles of gasP(total) is the total pressureWe can first calculate the total number of moles of gas in the mixture:n(total) = n(Ne) + n(Ar)where:n(Ne) is the number of moles of Nen(Ar) is the number of moles of ArWe can then calculate the number of moles of each gas using the following formula:n = m / Mwhere:n is the number of molesm is the massM is the molar massThe molar mass of Ne is 20.18 g/mol and the molar mass of Ar is 39.95 g/mol. Therefore, the number of moles of Ne and Ar are:n(Ne) = 10.0 g / 20.18 g/mol = 0.496 moln(Ar) = 10.0 g / 39.95 g/mol = 0.250 molTherefore, the total number of moles of gas in the mixture is:n(total) = 0.496 mol + 0.250 mol = 0.746 molFinally, we can calculate the partial pressure of Ne:P(Ne) = (0.496 mol / 0.746 mol) * 1.6 atm = 1.1 atmTherefore, the partial pressure of Ne is 1.1 atm.

Question 19

Using the graph below,determine the gas that has the highest density at STP.

Accepted Answer

The answer of Using the graph below,determine the gas that...

Question 20

Using the graph below,determine the gas that has the lowest density at STP.

Accepted Answer

The answer of Using the graph below,determine the gas that...

Quiz 5: Gases

Define hypoxia.

A compound is found to be 30.45% N and 69.55% O by mass.If 1.63 g of this compound occupy 389 mL at 0.00°C and 775 mm Hg,what is the molecular formula of the compound?

A syringe contains 589 mL of CO at 325 K and 1.2 atm pressure.A second syringe contains 473 mL of N₂ at 298 K and 2.6 atm.What is the final pressure if the contents of these two syringes are injected into a 1.00 L container at 0.00°C?

What pressure would a gas mixture in a 10.0 L tank exert if it were composed of 48.5 g He and 94.6 g CO₂ at 398 K?

Determine the volume of O₂ (at STP) formed when 50.0 g of KClO₃ decomposes according to the following reaction.The molar mass for KClO₃is 122.55 g/mol. 2 KClO₃(s) → 2 KCl(s) + 3 O₂(g)

To avoid oxygen toxicity and nitrogen narcosis,deep-sea divers breathe a specialized mixture of gas.Describe the mixture.

Determine the theoretical yield and the percent yield if 21.8 g of K₂CO₃ is produced from reacting 27.9 g KO₂ with 29.0 L of CO₂ (at STP) .The molar mass of KO₂ = 71.10 g/mol and K₂CO₃= 138.21 g/mol. 4 KO₂(s) + 2 CO₂(g) → 2 K₂CO₃(s) + 3 O₂(g)

A 0.334 g sample of an unknown halogen occupies 109 mL at 398 K and 1.41 atm.What is the identity of the halogen?

Determine the density of NH₃ gas at 435 K and 1.00 atm.

How many molecules of XeF₆are formed from 12.9 L of F₂ (at 298 K and 2.60 atm) according to the following reaction? Assume that there is excess Xe. Xe(g) + 3 F₂(g) → XeF₆(g)

Determine the volume of H₂S (at 375 K and 1.20 atm) needed to produce 55.0 g of S.Assume that there is excess SO₂ present. 2 H₂S(g) + SO₂(g) → 3 S(s) + 2 H₂O(g)

What is the volume of 5.60 g of O₂ at 7.78 atm and 415 K?

Define vapor pressure.

The total pressure of a gas mixture is the sum of the partial pressure of its components is known as

A 0.465 g sample of an unknown compound occupies 245 mL at 298 K and 1.22 atm.What is the molar mass of the unknown compound?

What mass of NO₂ is contained in a 13.0 L tank at 4.58 atm and 385 K?

A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm.What is the partial pressure of Ne?

Using the graph below,determine the gas that has the highest density at STP.

Using the graph below,determine the gas that has the lowest density at STP.

Quiz 5: Gases

Define hypoxia.

A compound is found to be 30.45% N and 69.55% O by mass.If 1.63 g of this compound occupy 389 mL at 0.00°C and 775 mm Hg,what is the molecular formula of the compound?

A syringe contains 589 mL of CO at 325 K and 1.2 atm pressure.A second syringe contains 473 mL of N2 at 298 K and 2.6 atm.What is the final pressure if the contents of these two syringes are injected into a 1.00 L container at 0.00°C?

What pressure would a gas mixture in a 10.0 L tank exert if it were composed of 48.5 g He and 94.6 g CO2 at 398 K?

Determine the volume of O2 (at STP) formed when 50.0 g of KClO3 decomposes according to the following reaction.The molar mass for KClO3 is 122.55 g/mol. 2 KClO3(s) → 2 KCl(s) + 3 O2(g)

To avoid oxygen toxicity and nitrogen narcosis,deep-sea divers breathe a specialized mixture of gas.Describe the mixture.

Determine the theoretical yield and the percent yield if 21.8 g of K2CO3 is produced from reacting 27.9 g KO2 with 29.0 L of CO2 (at STP) .The molar mass of KO2 = 71.10 g/mol and K2CO3 = 138.21 g/mol. 4 KO2(s) + 2 CO2(g) → 2 K2CO3(s) + 3 O2(g)

A 0.334 g sample of an unknown halogen occupies 109 mL at 398 K and 1.41 atm.What is the identity of the halogen?

Determine the density of NH3 gas at 435 K and 1.00 atm.

How many molecules of XeF6 are formed from 12.9 L of F2 (at 298 K and 2.60 atm) according to the following reaction? Assume that there is excess Xe. Xe(g) + 3 F2(g) → XeF6(g)

Determine the volume of H2S (at 375 K and 1.20 atm) needed to produce 55.0 g of S.Assume that there is excess SO2 present. 2 H2S(g) + SO2(g) → 3 S(s) + 2 H2O(g)

What is the volume of 5.60 g of O2 at 7.78 atm and 415 K?

Define vapor pressure.

The total pressure of a gas mixture is the sum of the partial pressure of its components is known as

A 0.465 g sample of an unknown compound occupies 245 mL at 298 K and 1.22 atm.What is the molar mass of the unknown compound?

What mass of NO2 is contained in a 13.0 L tank at 4.58 atm and 385 K?

A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm.What is the partial pressure of Ne?

Using the graph below,determine the gas that has the highest density at STP.

Using the graph below,determine the gas that has the lowest density at STP.

A syringe contains 589 mL of CO at 325 K and 1.2 atm pressure.A second syringe contains 473 mL of N₂ at 298 K and 2.6 atm.What is the final pressure if the contents of these two syringes are injected into a 1.00 L container at 0.00°C?

What pressure would a gas mixture in a 10.0 L tank exert if it were composed of 48.5 g He and 94.6 g CO₂ at 398 K?

Determine the volume of O₂ (at STP) formed when 50.0 g of KClO₃ decomposes according to the following reaction.The molar mass for KClO₃is 122.55 g/mol. 2 KClO₃(s) → 2 KCl(s) + 3 O₂(g)

Determine the theoretical yield and the percent yield if 21.8 g of K₂CO₃ is produced from reacting 27.9 g KO₂ with 29.0 L of CO₂ (at STP) .The molar mass of KO₂ = 71.10 g/mol and K₂CO₃= 138.21 g/mol. 4 KO₂(s) + 2 CO₂(g) → 2 K₂CO₃(s) + 3 O₂(g)

Determine the density of NH₃ gas at 435 K and 1.00 atm.

How many molecules of XeF₆are formed from 12.9 L of F₂ (at 298 K and 2.60 atm) according to the following reaction? Assume that there is excess Xe. Xe(g) + 3 F₂(g) → XeF₆(g)

Determine the volume of H₂S (at 375 K and 1.20 atm) needed to produce 55.0 g of S.Assume that there is excess SO₂ present. 2 H₂S(g) + SO₂(g) → 3 S(s) + 2 H₂O(g)

What is the volume of 5.60 g of O₂ at 7.78 atm and 415 K?

What mass of NO₂ is contained in a 13.0 L tank at 4.58 atm and 385 K?