October 22: Aspirin Titration Experiment

Description:

A titration is used to find the amount of a substance through causing a reaction that changes the pH. An indicator is used to show when the reaction is complete.

In this experiment, the aspirin(acertylsalic acid) reacted with  sodium hydroxide. Neutralization produces water and a salt, therefore:

           Equation- C₉H₈O₄(aq) + NaOH(aq) -> C₉H₇O₄Na(s) + H₂O

Data Processing:

We can use this equation to find out the number of moles of the aspirin in the titration through finding the number of moles of NaOH needed to react with the aspirin and how many moles there are present. The second part of the data processing is to find how much aspirin you get from each tablet, spending the same amount of money. First, you have to find the price of one tablet. Once you do that, you have to scale your values to the same price. Then you can compare the two amounts of aspirin to see which one contains more.

October 9: Making a Standard Solution

Making a Standard Solution

Description: 

A standard solution is a solution who's concentration you know exactly. Standard solutions are often used by chemists and it is essential that they are made accurately. This lab went through the basics of making a standard solution with Sodium Hydrogen sulfate.

Standard Procedure for Making a Standard Solution:

- Calculate the mass of the solvent that is required to make the required amount of the wanted solution.

              e.g. To make 250cm^(2) of 0.100 mol dm^(-3) we needed approximately 3.45g

-Weigh out the solvent, and pour it into a beaker. Weigh the weighing boat after you have poured out the solvent and subtract the traces from the previous recorded amount

              e.g. Our final weight ended up being approximately 3.44

-Add a portion of distilled water to the beaker and stir until dissolved entirely

              e.g. We used 100cm^(3)

-Pour the contents of the beaker into a volumetric flask using a funnel, and rinse the funnel and the beaker with distilled water. Fill the volumetric flask to the graduated mark.

              e.g. For our experiment we used a 250cm^(3) volumetric beaker. 

The same steps are illustrated in the picture below:

October 20: Back-Titrations

October 20 : Back-Titrations

  *   *   *   *   *   *   *   *   *  

Today in class, we conducted an experiment which dealt with new concepts of stoichiometry: titrations. We were able to choose between two alternatives (eggshells or aspirin analysis). The one that we decided to work on was the eggshell experiment. 

Brief Description:

The major component of eggshells is CaCO3, known as calcium carbonate. We had to react ≈ 0.40g of CaCo3 with excess 1M of HCl, and when the reaction was complete, we also had to titrate the unreacted HCl with 1M of NaOH (acid-base titration). With further calculations, we were able to figure out the amount of calcium carbonate present and the amount of HCl that actually reacted with it.

Objectives:

  1. Determine the amount of CaCo3 present in the eggshell used

  2. Relate experimental titration measurements to a balanced chemical equation

  3. Infer a conclusion from experimental data

  4. Apply reaction-stoichiometry concepts

Procedure:

  1. Prepare your data table.

  2. Wash the pipette with the 1.0M HCl supplied.

  3. Wash the burette with the 1.00M NaOH supplied and fill it up.

  4. Weigh accurately 0.40g (this means weigh around 0.4g, but record the actual reading to 2 decimal places) of the powder and put it into the conical flask. 

  5. Add a few drops of ethanol to help the acid react with the calcium carbonate.

  6. Pipette onto the powder, 10cm³ of HCl and add a few drops of phenolphthalein indicator. Swirl the flask to assist reaction.

  7. Now, titrate the excess acid in the flask against the NaOH from the burette until a faint pink colour appears.

  8. Repeat steps 5-7 as necessary.

Collected Data:

  * Accurate mass of shells 

  * Volume of NaOH before and after titrating

Calculations!:

  1. First, find the change of volume of NaOH. You can figure this out by subtracting the before-volume from the after-volume. This will give you the amount of NaOH used to neutralize the unreacted acid.

  2. Then, find the number of moles of the initial HCL used in this experiment. 

  3. Next, looking at your titration data, work out the number of moles for NaOH; this will also give you the # of moles of HCl (molar ratio used).

  4. Calculate how much of the acid has reacted!

  5. Convert out of moles :) Good job, you've solved the amount of the reacted HCl!

  6. Now, we have to solve how much of CaCO3 is present. To start off, look at the molar ratio of your balanced equation. 

                               2HCl(aq) + CaCO3(s) → CaCl2(aq) + CO2(g) + H2O(l) + HCl(aq)

We can see that the ratio is 2:1 (HCl : CaCO3). And since we already know the # moles of the reacted HCl, you can easily divide that value by 2 to get the # moles of calcium carbonate. Again, convert out of moles and answer the question correctly!

  

Conclusion:
  
See if you can adress the problem that the farmer was concerned about, possibly by calculating the percent of CaCO3 present in the actual eggshell. And also remember to state the limitations and certain errors (random / systematic)!


Assignments:

Finish this lab and STUDY FOR THE TEST! GOOD LUCK!

October 03: Behaviour of Gases

Grade 11 Chemistry Blog
October 03: Behaviour of Gases
Assignments due: How Big is One Mole of Gas?

During class, we explored the behaviour of gases when the temperature and pressure were changed. We did this through Experiment A and Experiment B. The relationship between the volume of gas and the temperature/pressure was recorded in Logger Pro and could be analysed using the graphs.

Experiment A (Pressure vs Volume):
In experiment A, we explored the relationship between pressure and volume. We did so by changing the volume of a gas and recording its pressure. The relationship found was that pressure is inversely proportional to volume.
We learnt about this in Boyle's Law. We can state Boyle's Law as the following:


We found this relationship by measuring the volume and pressure of a gas in a syringe when the volume is changed. The Logger Pro data we found looked like this:

However, if we changed the values of volume to be a reciprocal on Logger Pro, we were able to analyse a linear graph and prove Boyle's Law. The processed graph would look like this:


Experiment B (Pressure vs Temperature):
Our second experiment was to find the relationship between temperature and pressure. We did so by changing the temperature of a gas and recording its pressure. The relationship found was that pressure is proportional to temperature.
This relationship is described in Gay-Lussac's Law which can be expressed as:

The relationship was found by heating water and placing a flask of gas in different temperatures. The pressures recorded in the Logger Pro was expressed in a linear graph and proved Gay-Lussac's Law. The Logger Pro graph looked like this:

October 7: Determining the Molar Mass of Butane 

Determining the Molar Mass of Butane

Today, depending how far you previous got last class on friday, you either conducted the lab Determining the Molar Mass of Butane, or if you had already conducted this lab, you finished processing the calculations for the data/observations and Post Lab Questions. 

The objective of the lab: 

• To successfully collect and store butane gas from a cigarette lighter.
• To use the Ideal Gas Law (and Dalton’s Law of Partial Pressures) in calculating the number of moles of butane being stored.
• To use the mass and the number of moles of butane being stored in calculating the molar mass of butane.
• To compare the experimental molar mass of butane with the theoretical mass (based on its formula, which will be provided) and calculate a percent error.

For those who conducted the lab today, followed the procedure on there Butane Lab sheet which involved -

• Obtain a 2 or 3 dp balance and make sure it is properly set to zero.
• Place a clean, dry cigarette lighter on the triple beam balance and record its mass under Initial Mass of Cigarette Lighter in the Data and Observations section.  Remember to record the mass in agreement with the uncertainty of the device (if you can). 
• Add water to a trough (at least 2/3 full).  
• Fill a 500cm³ or 1dm³ measuring cylinder with water to the very brim.  Place a glass plate over the mouth of the bottle, then turn the cylinder upside down and place it in the trough.  When the mouth underwater, you may remove the glass plate.  Be sure that there are no bubbles in the bottle.  If the bottle has air bubbles trapped in it, you must try again to completely fill the bottle with water.
• Lift the bottle up slightly (but keep its mouth below the surface of the water) and hold the cigarette lighter beneath the mouth.  Press the button on the lighter so that bubbles of butane rise into the inverted bottle.  Be careful that no bubbles of butane miss the mouth of the bottle.
• As the cylinder fills with butane, the level of the water will fall.  Continue to collect butane gas until at least 200cm³ have been collected.
• When you have collected a measurable amount of butane, release the button on the lighter and remove it from the water.  Dry it with a paper towel and hair dryer and record its mass again under Final Mass of Cigarette Lighter in the Data and Observations section.  The second mass should be LESS than the first mass you recorded in step A.  If the second mass is higher than or the same as the first mass, be sure that it is thoroughly dry (including inside the metal casing around the gas outlet) and mass it again.  If the mass is STILL greater than or the same as the first mass, you should repeat the procedure from the beginning.
• Read the volume of the butane you collected from the side of the cylinder.  Record this volume under Volume of Butane Collected in the Data and Observations section.  When you have recorded the volume, you may set the bottle down inside the trough.
• Use a thermometer to record the temperature of the water in the trough.  You may assume that the gas temperature is the same as the temperature of the water.  Record the temperature to the correct number of decimal places under Temperature of Gas Collected in the Data and Observations section. 
• Use a weather website to identify the atmospheric pressure in the room.  This information will be recorded on the board in mmHg.  Record this pressure in your Data and Observations section exactly as it appears on the board under the heading Total Pressure of Gas in Cylinder.
• Refer to the graph on the next page to determine the vapor pressure of H₂O at the temperature you recorded in Step I.  Record the Vapor Pressure of Water in your Data and Observations section.
• Perform all calculations and conversions in the Calculations section.  Don’t forget to calculate a percent error when you have finished everything else.

Once everyone had finished this procedure, they continued to work on the calculations starting with the Data and observations worksheet for the lab which included the following:

Initial Mass of Cigarette Lighter:  _________________________ grams

Final Mass of Cigarette Lighter:  _________________________ grams

Volume of Gas Collected in mL:  _________________________ mL



Temperature of Gas Collected in ºC:  _____________________ºC



Atmospheric Pressure in mmHg*:  __________________ mmHg

Vapor Pressure of Water at Measured Temperature:  _________________ mmHg

This Process also involved the need of this graph below to find out some of the values. 

Once all the Observations and data was processed students moved on to the 5 Post Lab Questions which applied the same concepts as the above valuations.  The Questions are as follows: 

1.Identify at least one potential source of error in the experimental procedure that may lead to a loss of accuracy.

2.A gas sample is collected over water at 20.0ºC.  The volume of the gas collected is 45.0 mL and the atmospheric pressure is 771 mmHg.  How many moles of gas were collected?  (Remember to take into account the vapor pressure of water.  Consult the graph on Page 3 of this lab.)

3.The mass of the gas sample described in question 2 is found to be 0.0371 grams.  What is the molar mass of this gas?

4.It is known that the gas described in questions 2 and 3 is an element from the periodic table.  What is the probable identity of this mystery gas?

5.What volume will 0.010 grams of n-pentane, C₅H₁₂, occupy when stored under 12.0 atm of pressure and 21.0ºC?

In both worksheets students used the concept of the Ideal Gas Law:  PV=nRt. 
Also students applied the relations between the Molar Volume of a Gas, converting Temperatures and Pressures into the correct units. 

•  Which can be found on your (Gas Law Practice Problems 35 mins) handout. 

Homework: Finish Butane lab work, Watch/take notes on video 1.33 and prepare for the quiz on section 1.32 on this thursday. Continue with all worksheets. 

October 01: Behavior of Gases

After starting off with a quiz from syllabus item 1.31: "Stochiometry reaction masses  and volume (1)", we used Gay-Lussac and Boyle's law to calculate the behavior of gases.

Experiment A

Boyle's law 
Pressure is inversely proportional to volume.
We find this through the equation
PV=k
where k is a constant
We used this in a pressure-volume relationship for gases lab
By measuring the pressure in a syringe when we decreased the volume in the syringe. The pressure was measured by using a pressure sensor which we then got plotted into LoggerPro on a computer and choosing which volume we would have.


As the relationship was inversely proportional when we dotted the data we got a hyperbola.
There were several tasks concerning this lab, concidering the inversely proportional relationship, where if the volume is doubled what would happen to the pressure, and if the volume was halved, what would then happen to the pressure.
 For example when the volume goes from 10 to 5 mL, the pressure will double.

Experiment B

Gay-Lussac's law 

P/T=k
Gay-Lussac's law states that pressure is proportional to temperature.
Even though we didn't finish the second experiment, we still did the first part of it.
The experiment was about a pressure-temperature relationship and so when the temperature increased so would the pressure. This was proven by putting an 125 mL Erlenmeyer flask into 800 mL of water with different temperatures with the pressure sensor attached so , and then use LoggerPro to collect data and get a linear fit. The temperatures varied from ca. 48 C to 18 C.