# Lab : Gases, Pressure and Stoichiometry

Add 0.5g of sodium bicarbonate, NaHCO3 into a flask. Record the initial pressure and temperature in the flask. Add 10 mL of water to the flask to dissolve the NaHCO3 then record the pressure as the initial pressure for the experiment. Record the gas volume as the initial volume. Add 1 mL of 1M HCl, allow the temperature in the flask to return to room temperature, and then record the pressure in the flask. Continue to add 1 mL increments of HCl solution. (observe whether a reaction has occurred).

These are the initial conditions of the closed flask with just the sodium carbonate inside. Assume that the sodium bicarbonate powder takes up a negligible volume in the flask.

(a) Initial pressure (atm): 1.072 atm

(b) Initial gas volume (mL): 139.72 mL

(c) Initial temperature (C): 21C

These are the final conditions of the closed flask:

(a) Final pressure (atm): 2.405 atm

(b) Final gas volume (mL): 131.92mL

(c) Final temperature (C): 21 C

1. Use the initial conditions and the ideal gas law to calculate the number of moles of air molecules (N2, O2, etc.) trapped in the flask. The gas constant, R = 0.082057 (L*atm)/(K*mole), so make sure to convert your temperature value to Kelvin.

2. The final conditions represent the air and CO2 from the reaction held in a smaller volume than at the start. Calculate the total moles of air and CO2 molecules from the ideal gas law, using the final conditions of pressure, volume and temperature. Record the total moles here.

3. Subtract the moles of air (from the initial conditions) from the total moles of gas at the end to find the moles of CO2 produced by the reaction.Record the moles of CO2 gas produced by the reaction:

4. You now have enough information to balance the equation: NaHCO3(s) + HCl(aq) --> CO2(g) + H2O(l) + NaCl(aq)

Calculate the following:

(a) Moles of NaHCO3 added to the flask (MW=84.007 g/mol):

(b) Moles of CO2 produced

(c) The ratio of (moles of CO2) divided by (moles of NaHCO3)

(d) What is the nearest whole number ratio for your experimentally-determined ratio? For example, if the experimental ratio is 1.75, then you would use a whole number ratio of 2 to 1. If the experimental ratio is 3.2, then you would use a whole number ratio of 3 to 1.

(e) Use the whole number ratio to balance the equation for the reaction. If your whole number ration is 2 to 1, then place a 2 next to the CO2 and a 1 next to the NaHCO3. Continue from there to balance the number of atoms on each side of the equation. Write the balanced equation here:

(f) Calculate the experimental error from the difference between the experimental ratio and the whole number ratio as follows:

% error = |real ratio - whole number ratio| / (whole number ratio) * 100%

5. (a) Could you boil off the liquid at the end of the experiment and weigh the NaCl left in the flask in order to balance the equation?

(b) Air usually contains water vapor. The air trapped in the flask was dry air with no water vapor. How would the presence of water vapor affect your calculations?

(c) Carbon dioxide is slightly soluble in water and its solubility increases with increasing pressure. How would the fact that some of the carbon dioxide produced by the reaction is dissolved in the aqueous solution affect your calculations?

https://brainmass.com/chemistry/gas-stoichiometry/lab-gases-pressure-and-stoichiometry-407259

#### Solution Preview

These are the initial conditions of the closed flask with just the sodium carbonate inside. Assume that the sodium bicarbonate powder takes up a negligible volume in the flask.

(a) Initial pressure (atm): 1.072 atm

(b) Initial gas volume (mL): 139.72 mL

(c) Initial temperature (C): 21C

These are the final conditions of the closed flask:

(a) Final pressure (atm): 2.405 atm

(b) Final gas volume (mL): 131.92mL

(c) Final temperature (C): 21 C

1. Use the initial conditions and the ideal gas law to calculate the number of moles of air molecules (N2, O2, etc.) trapped in the flask. The gas constant, R = 0.082057 (L*atm)/(K*mole), so make sure to convert your temperature value to Kelvin.

Use PV = nRT (all units must match units of R)

R = 0.082057 (L*atm)/(K*mole)

P = 1.072 atm

V = 139.72 mL x (1 L/1000 mL) = 0.13972 L

T = 21 C = 21 + 273 = 294 K

n = PV/RT = (1.072 atm x 0.13972 L)/(0.082057 (L*atm)/(K*mole) x 294 K) = 0.00621 moles

2. The final conditions represent the air and CO2 from the reaction held in a smaller volume than at the start. Calculate the total moles of air and ...

#### Solution Summary

Stoichiometric calculations relating to a reaction that produces a gas are illustrated. Results are obtained from the PV = nRT reaction. Mole ratios of reactants and products are also investigated.

Molar Volume of an Ideal Gas Lab

Part 1

The following are the givens from the experiment:

Initial Experiment with 10mL- 6 M HCl and 0.25g Zinc

Initial Teperature = 21.5C

Initial Pressure = 1.00atm

After adding 0.25g Zinc - The solution reacted

Pressue = 1.69atm

Temperature = 27.3C

Volume in Syrenge = 92.2mL

Volume in Flask = 9.9mL

Part 2

Repeated Experiment with 10mL- 6 M HCl and 0.5g of Zinc

Initial Temperature = 21.5C

Initial Pressure = 1.00atm

After adding 0.5g Zinc -The solution reacted

Pressure = 2.41atm

Temperature =33.4C

Volume in Syrenge = 184.5mL

Volume in Flask = 9.8mL

The following answers are what I have so far: Need to know if A,B,C,D,E, & F are correct for 1 & 2 if not please correct.

1. Record and calculate the following for the first half of the experiment:

a. Mass of zinc used (g) = 0.25g

b. Volume of 6M HCl used (mL) = 10mL

c. Molecular weight of zinc (g/mol) = 65.409 g/mol

d. Moles of zinc reacted = 0.25 g / 65.41 g/mol= 0.00382 moles

e. Moles of hydrogen produced = 0.00382 moles

f. Molar volume of the ideal hydrogen gas at room temperature (Volume/moles), expressed as L/mol at X degrees C and a pressure of 1 atmosphere = 22.4 L/mole * 0.00382 moles = 0.0856 L or 85.6mL

2. For the second part of the experiment, everything was the same except that twice as much Zn was used. Record and calculate the following for this second reaction:

a. Mass of zinc used (g) = 0.50

b. Volume of 6M HCl used (mL) = 10mL

c. Molecular weight of zinc (g/mol) = 65.409 g/mol

d. Moles of zinc reacted = 0.50 g / 65.41 g/mol=0.00764

e. Moles of hydrogen produced = 0.00764 moles

f. Molar volume of the ideal hydrogen gas at room temperature (Volume/moles), expressed as L/mol at X degrees C and a pressure of 1 atmosphere = 22.4 L/mole * 0.00764 moles = 0.1711 L

Also need the following answered from the above information: Numbers 3, 4, & 5

3. Compare the molar volumes obtained in the two parts of the experiment. What difference did using twice the amount of Zn in the second part?

4. Which is the limiting reactant - zinc or HCl? I think HCl

5. Compare the experimental value for the molar volume at 21C with the value listed in the Background section of the lab manual. Calculate the experimental error according to: % error = | experimental molar volume - listed molar volume | / (listed molar volume) * 100%

Background

Robert Boyle found that, under isothermal conditions (constant temperature), the pressure of a gas is inversely proportional to its volume. This also means that the product of the pressure and volume is a constant. The equations that describe this relationship are

P × V = constant

and

P1 × V1 = P2 × V2

Jacques Charles found that the volume of a gas under isobaric conditions (constant pressure) is directly proportional to its temperature. As the temperature increases so does the volume that the gas occupies. The equations that describe this relationship are:

V ÷ T= constant

and

V1 ÷ T1 = V2 ÷ T2

It is possible to combine Boyle's law and Charles' law:

P1 × V1 ÷ T1 = P2 × V2 ÷ T2

The Ideal Gas Law includes the dependence of the volume on the number of gas moles, and an experimentally determined proportionality constant, called the Universal Gas Constant, which is designated by R.

The mathematical formulation of the law is written as follows:

(P × V) ÷ (n × T) = R

or in the more familiar form:

P × V = n × R × T

The value for the Universal Gas Constant, R, is 8.314472 J/mol*K (Joules per mole per degree Kelvin).

One important result of the Ideal Gas Law is that under conditions of constant pressure and temperature, one mole of any gas will always occupy the same volume.

The volume that one mole of an ideal gas occupies when held at a specific temperature and pressure is referred to as a "molar volume". For example, at one atmosphere, the molar volume of any ideal gas is 22.414 L/mol at 0 °C and 24.137 L/mol at 21°C. Most gases follow the Ideal Gas Law closely at atmospheric pressure and room temperature.

In this experiment you will determine the molar volume of hydrogen gas and compare it with the value predicted by the Ideal Gas Law. Hydrogen gas is produced in the following reaction between zinc and hydrochloric acid:

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

One can calculate the molar volume of the gas at room temperature and pressure according to the equation:

Molar volume (L/mol) = (measured volume) ÷ (number of moles)

To complete this lab you will need to:

Calculate the number of moles of H2 produced. (Hint: For every mole of zinc used in the reaction, one mole of H2 is produced.)

Measure the volume of H2 gas produced.