I really need help answering these questions.

A sample of CO2 gas with a mass of 0.30 g is placed in a 250 mL container at 400 K. What is the pressure in atmospheres exerted by this gas?

Before you can put the numbers into an equation and solve for the volume, change the mass to moles by using the following equation:

1)mass*molar mass

2)mass/molar mass

3)mass*(6.02 x 1023)mass/

4)mass(6.02 x 1023).

If you calculate the molar mass of CO2, what do your get? g/mol (round the atomic masses to hundredths place and your answer should also be rounded to the hundredths place)

Now, use the equation you chose above to calculate the moles of CO2. mol(round to the ten thousandths place)

Almost ready to calculate our answer, but first we need to convert the mL to L.
250 mL x 1)1000L/mL 2)1L/1000mL 3)100L/ mL 4)1L/100mL = our answer in L.

After plugging your data into the correct ideal gas law equation, you get LatmpsimL (round to the hundredths place).

You want to use PV = nRT

It appears that someone has set up the problem to work it step by step. Can you not follow those steps? Has someone tried to help and now you're stuck. Are these instructions from your teacher? I can help you through it but I need to know the answers to the above questions before I start. Thanks for using Jiskha.

Yes the instructions were from my teacher.

Well, well, well, looks like you've got some chemistry questions there! Don't worry, I'm here to help you with a sprinkle of humor. Let's break it down, shall we?

First, to convert the mass of CO2 to moles, we need to divide the mass by the molar mass of CO2. It's like finding out how many cookies are in a jar by dividing the total weight of cookies by the weight of one cookie. So go ahead and choose the equation that does that for you!

Now, let's calculate the molar mass of CO2. Remember, we need to round the atomic masses to the hundredths place. Though I must say, carbon and oxygen must be quite fashionable atoms to wear decimal points on their outfits.

Next, it's time to calculate the moles of CO2 using the equation you picked earlier. Think of it as figuring out how many cookies are in the jar by dividing the total weight of cookies by the weight of one cookie.

But wait, there's more! We need to convert milliliters to liters, just like converting inches to feet. Which equation will help us make that conversion? Choose wisely!

Finally, plug in all the data you've gathered into the correct ideal gas law equation and see what our answer is! I won't make any more puns here because I don't want to gas you out, but remember to round your answer to the hundredths place.

Good luck, and may the chemistry be with you!

To answer the first question, we need to determine the moles of CO2 gas using the given mass and molar mass. The correct equation to convert mass to moles is:

2) mass / molar mass

To calculate the molar mass of CO2, we need to look up the atomic masses of carbon (C) and oxygen (O) and multiply them by the respective subscripts in the chemical formula CO2.

The atomic mass of carbon (C) is approximately 12.01 g/mol, and the atomic mass of oxygen (O) is roughly 16.00 g/mol. Therefore, the molar mass of CO2 is:

(1 * 12.01 g/mol) + (2 * 16.00 g/mol) = 44.01 g/mol (rounded to the hundredths place)

Using this molar mass, we can now calculate the moles of CO2:

moles of CO2 = (mass of CO2) / (molar mass of CO2)

moles of CO2 = 0.30 g / 44.01 g/mol ≈ 0.00682 mol (rounded to the ten thousandths place)

Next, we need to convert the given volume of 250 mL to liters. The correct equation to convert mL to L is:

2) 1 L / 1000 mL

Thus, the conversion is:

250 mL * (1 L / 1000 mL) = 0.25 L

Now, we can use the ideal gas law equation to find the pressure exerted by the gas. The ideal gas law equation is:

P * V = n * R * T

where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant (approximately 0.0821 L·atm/(mol·K)), and T is the temperature in Kelvin.

Rearranging the equation to solve for P, we get:

P = (n * R * T) / V

Plugging in the values we have:

P = (0.00682 mol * 0.0821 L·atm/(mol·K) * 400 K) / 0.25 L ≈ 8.74 atm (rounded to the hundredths place)

Therefore, the pressure exerted by the CO2 gas is approximately 8.74 atmospheres.

Are you sure you copied the instructions correctly? Some of them don't make sense to me.

I have copied the question to this space.
A sample of CO2 gas with a mass of 0.30 g is placed in a 250 mL container at 400 K. What is the pressure in atmospheres exerted by this gas?

Here is how you do the problem.
Use PV = nRT, the general gas equation.
P is pressure in atmospheres. Solve for this. That is the only unknown in the equation. The answer will be in the unit atmospheres.

V = volume in liters. The problem gives 250 mL so change that to 0.250 L.

n = number of mols of gas. The problem gives 0.30 grams. mols = grams/molar mass = 0.30/44 = ?? (I have estimated the molar mass of CO2. You should confirm that number.)

R is the universal gas constant, which for this problem is 0.08206 L*atm/mol*K.

T is the temperature in Kelvin which the problem lists as 400 K.
Now, just plug those numbers in and solve for P. I'll be happy to check your answer.