Q1. Consider the following data table about atmospheric pressure units. (a) Which units are shown, and how do they relate?

Background

Topic: Units of Pressure

This question tests your understanding of the different units used to measure atmospheric pressure and how to convert between them.

Key Terms and Formulas:

• Atmosphere (atm)

• Millimeters of mercury (mm Hg or torr)

• Pascals (Pa)

• Pressure conversion factors (e.g., )

Step-by-Step Guidance

1. Identify each unit listed in the table and write out what it stands for (e.g., atm, mm Hg, Pa).

2. Recall the standard conversion factors between these units.

3. Explain how these units are related using the conversion factors.

Try solving on your own before revealing the answer!

Q2. Explain briefly why the pressure decreases as the altitude increases.

Background

Topic: Atmospheric Pressure and Altitude

This question is about the relationship between atmospheric pressure and altitude, and why pressure changes as you go higher above Earth's surface.

Key Terms:

• Atmospheric pressure

• Altitude

• Air density

Step-by-Step Guidance

1. Think about what causes atmospheric pressure (the weight of air above a given point).

2. Consider what happens to the amount of air above you as you go higher in altitude.

3. Relate the decrease in air molecules above you to the decrease in pressure.

Try solving on your own before revealing the answer!

Q3. At a low place on Earth (such as below sea level), would you expect the atmospheric pressure to be more or less than at sea level? Why?

Background

Topic: Atmospheric Pressure and Elevation

This question asks you to apply your understanding of how atmospheric pressure changes with elevation, specifically below sea level.

Key Terms:

• Atmospheric pressure

• Elevation (below sea level)

Step-by-Step Guidance

1. Recall how atmospheric pressure changes with altitude (from previous question).

2. Think about whether being below sea level means there is more or less air above you compared to sea level.

3. Predict whether the pressure would be higher or lower and explain why.

Try solving on your own before revealing the answer!

Q4. Describe with words or pictures how temperature and pressure are related. You must address what is happening at the particle level.

Background

Topic: Kinetic Molecular Theory (KMT)

This question tests your understanding of how temperature affects the motion of gas particles and, in turn, the pressure they exert.

Key Terms:

• Kinetic energy

• Gas particles

• Pressure

• Temperature

Step-by-Step Guidance

1. Recall that temperature is a measure of the average kinetic energy of gas particles.

2. Think about what happens to the speed of particles as temperature increases.

3. Explain how faster-moving particles collide with the walls of the container more often and with greater force, affecting pressure.

Try solving on your own before revealing the answer!

Q5. Convert a pressure of 0.600 atm to the following units: (a) mm Hg, (b) Pa

Background

Topic: Pressure Unit Conversions

This question tests your ability to convert between different units of pressure using standard conversion factors.

Key Formulas:

Step-by-Step Guidance

1. Write down the given pressure in atm.

2. Multiply by the appropriate conversion factor to get mm Hg.

3. Multiply by the appropriate conversion factor to get Pa.

Try solving on your own before revealing the answer!

Q6. A 250 mL sample of gas is collected at 57°C. What volume will the gas sample occupy at 25°C if the pressure is constant?

Background

Topic: Gas Laws (Charles's Law)

This question tests your ability to use Charles's Law to relate the volume and temperature of a gas at constant pressure.

Key Formula:

• Charles's Law:

Where:

• = initial volume

• = initial temperature (in Kelvin)

• = final volume

• = final temperature (in Kelvin)

Step-by-Step Guidance

1. Convert both temperatures from Celsius to Kelvin by adding 273.15.

2. Write down the initial and final values for volume and temperature.

3. Set up Charles's Law equation with the known values.

4. Rearrange the equation to solve for the unknown volume ().

Try solving on your own before revealing the answer!

Q7. Which of the following equations represents the oxidation of ammonia gas? (Choose the correct balanced equation.)

Background

Topic: Chemical Equations and Balancing

This question tests your ability to recognize and balance chemical equations, specifically for the oxidation of ammonia ().

Key Terms:

• Oxidation

• Ammonia ()

• Balanced chemical equation

Step-by-Step Guidance

1. Recall what oxidation means (loss of electrons, often involving reaction with ).

2. Identify the reactants and products in the oxidation of ammonia.

3. Check each equation for correct balancing of atoms on both sides.

Try solving on your own before revealing the answer!

Q8. A 1.0 L sample of gas at 430 K exerts a pressure of 2.0 atm. If the gas is compressed to 1.2 L at constant temperature, what is the new pressure?

Background

Topic: Boyle's Law

This question tests your ability to use Boyle's Law to relate the pressure and volume of a gas at constant temperature.

Key Formula:

• Boyle's Law:

Where:

• = initial pressure

• = initial volume

• = final pressure

• = final volume

Step-by-Step Guidance

1. Write down the known values for , , and .

2. Set up Boyle's Law equation with the known values.

3. Rearrange the equation to solve for the unknown pressure ().

Try solving on your own before revealing the answer!

Q9. Calculate the volume that contains 6.0 x 1023 chlorine molecules (Cl2) at STP. The calculation can be represented by the equation: . (Hint: 1 mole = 6.02 x 1023 molecules; STP = 1 atm, 273 K; R = 0.0821 L·atm/mol·K)

Background

Topic: Ideal Gas Law and Moles

This question tests your ability to use the ideal gas law to find the volume of a gas, given the number of molecules at standard temperature and pressure (STP).

Key Formula:

• Ideal Gas Law:

• Number of moles:

Where:

• = pressure (atm)

• = volume (L)

• = number of moles

• = ideal gas constant (0.0821 L·atm/mol·K)

• = temperature (K)

Step-by-Step Guidance

1. Calculate the number of moles of Cl2 using Avogadro's number.

2. Write down the values for , , , and at STP.

3. Set up the ideal gas law equation to solve for .

4. Plug in the values, but stop before calculating the final volume.

Try solving on your own before revealing the answer!

Q10. Which gas diffuses faster: O2 or CO2? Explain why.

Background

Topic: Graham's Law of Effusion/Diffusion

This question tests your understanding of how the rate of diffusion of a gas depends on its molar mass.

Key Formula:

• Graham's Law:

Where:

• and are the rates of diffusion of gases 1 and 2

• and are the molar masses of gases 1 and 2

Step-by-Step Guidance

1. Find the molar masses of O2 and CO2.

2. Set up Graham's Law to compare the rates of diffusion.

3. Predict which gas will diffuse faster based on the relationship between rate and molar mass.

Try solving on your own before revealing the answer!