Sciencemadness Discussion Board

Benchmarks in learning chemistry

Atrum - 16-2-2015 at 20:57

It has been roughly a year since I started to focus on learning chemistry and I find my self randomly picking experiments to do in order to further my knowledge in the field.

I was wondering if someone can recommend a series of chemistry experiments in order of increasing difficulty.

Basically I am looking for progressive series of chemistry projects whether they be synthesis, analysis, or isolation set up in such a way that each builds off the experience of the last.


I know this may seem silly, but I am simply looking for a way to add some structure to my learning adventure with out actually taking a chemistry course. My current career does not leave me with adequate time or money to do so.

Thanks,
Atrum

Chemosynthesis - 16-2-2015 at 21:00

This probably isn't as helpful as you would like, but why not use a lab book? There are several online for free. Is it that you would like a slightly more condensed structure to the experimentation due to time concerns?

j_sum1 - 16-2-2015 at 21:07

Well, Chemistry knowledge and understanding builds on itself but there is no prescribed order. You will find here some people who have a bit of knowledge of inorganic but have little more than the elementary concepts of organic. Likewise there are those with a huge familiarity with the chemicals themselves and their properties but have little grasp of thermodynamics. If it was an entirely linear path then I think your request would be easy.

I would grab a textbook that covers the things you wish to learn and progress through it. A good text would have some experiments to attempt as well. Then you could frame your request differently: "I don't have any silver nitrate. Can anyone suggest a similar or better precipitation reaction than this one?" or whatever.

Not that I have looked for a while but there are some great free rsources on the SM library.

http://library.sciencemadness.org/library/index.html

[Edited on 17-2-2015 by j_sum1]

Atrum - 16-2-2015 at 21:37

Thank you for the replies.

I will take your advice and check out the library.
In addition to what is available on sciencemadness. Are there any lab manuals written in recent times that you can recommend as well?

MrHomeScientist - 17-2-2015 at 11:25

Robert Bruce Thompson's "The Illustrated Guide to Home Chemistry" is a really great book. It's full of information on lab equipment and techniques, and takes you through a lot of fun experiments. I believe several places sell kits that include all the chemicals you need for the experiments in the book, too. I highly recommend it!

Atrum - 21-2-2015 at 13:30

Thanks again.

aga - 21-2-2015 at 13:56

Make 98% Nitric Acid.

It is great distillation to do.

Slightly scary, but looks great.

aga - 21-2-2015 at 14:59

Your OP was a bit random.

Better to ask for suggestions about what to do next, and say what you have already done.

E.g. do you understand molecular weights and stoichiometry ?

Atrum - 21-2-2015 at 15:01

Quote: Originally posted by aga  
Make 98% Nitric Acid.

It is great distillation to do.

Slightly scary, but looks great.


I have a list of synthesis reactions I want to attempt. That is one of them.

Lately my focus has been on titration of acids and bases. I find I have trouble detecring the very faint colour of pink from the phenolphthalien making the end point hard to spot.

aga - 21-2-2015 at 15:05

"very faint colour of pink from the phenolphthalien"

Huh ?

Two possible causes :

1. You're adding the Titrator too fast
2. The Phenolphtalein is too old.

Atrum - 21-2-2015 at 15:32

Hmm. Well the phenolphthalien I have was purchased less than a year ago.

How fast is too fast when adding the titrant?

During my initial titration of sulfuric acid I find that if I add the titrant to a point that the pink colour is obvious then I get a molarity that isn't possible.

Unless I am mistaken, the molarity of sulfuric acid can't exceed 18 molar. The results I got showed ~19 molar.

Oscilllator - 22-2-2015 at 22:20

Quote: Originally posted by Atrum  
Hmm. Well the phenolphthalien I have was purchased less than a year ago.

How fast is too fast when adding the titrant?

During my initial titration of sulfuric acid I find that if I add the titrant to a point that the pink colour is obvious then I get a molarity that isn't possible.

Unless I am mistaken, the molarity of sulfuric acid can't exceed 18 molar. The results I got showed ~19 molar.

Towards the endpoint, you should be adding single drops at a time and waiting several seconds in between drops to watch for the phenolphthalien colour.
I think though that this is not the cause of you getting an impossible molarity, unless you are titrating ~98% sulfuric acid. More likely you have messed up the calculations (H2SO4 is diprotic!) or your chemicals are impure. What are you using as a titrant?

Atrum - 28-2-2015 at 01:03

Sorry for the late reply. My job tends to get in the way of my chemistry hobby.

I am using Na2CO3 (sodium carbonate). I am unsure as to the purity though. My source is called Super Washing Soda. The box states that it is 100% Sodium Carbonate.

I dried roughly 200g in an oven set to 230 F for about 2 hours.

I prepared a 1 molar solution of Na2CO3.

Sometimes I get a molarity of ~18.2 other times I get ~19.2; so the only thing I can think of is as you said the purity of my reagent may be the issue.

I know that H2SO4 is diprotic and Na2CO3 when mixed with water produces 2OH

My calculations are as follows:
Va = 5mL
Ca = X

Vb = 96.2mL
Cb = 1.0m

The coefficients for both are 2 which cancel so I have omitted them from the calculation.

(1m)(0.0962L) / 0.005L = X = 19.24m

I have gotten a range from ~18 to my highest which is the result I wrote above.


Sorry if my post isn't very well thought out, I am writing this after a 48 hour shift at work.

Thanks,
Atrum.
[Edited on 2-28-2015 by Atrum]

[Edited on 2-28-2015 by Atrum]

[Edited on 2-28-2015 by Atrum]

diddi - 28-2-2015 at 15:22

you need to be very suspicious of any results based on a standard solution that was sourced from a laundry box. you need to find a reliable standard. or standardise your Na2CO3.

Atrum - 28-2-2015 at 16:01

Thanks, I guess I should have known better. If I need a standard solution I had better order something I know is pure.

morganbw - 28-2-2015 at 18:51

@Atrum
Hey guy/gal, it looks like you are doing great first steps.
Keep working @ it.

Atrum - 28-2-2015 at 19:45

Thanks Morganbw.

Oscilllator - 1-3-2015 at 16:52

Quote: Originally posted by Atrum  
Thanks Morganbw.

I would suggest recrystallising your Na2CO3 at least twice, then drying it in your oven as high as it will go in a clean glass dish. This should get you to a reasonable level of purity.
One other thing that could be introducing error is your equipment. Are you using burettes, pippets and volumetric flasks in the proper manner to measure out all your liquids?

Atrum - 1-3-2015 at 17:12

Quote: Originally posted by Oscilllator  
Quote: Originally posted by Atrum  
Thanks Morganbw.

I would suggest recrystallising your Na2CO3 at least twice, then drying it in your oven as high as it will go in a clean glass dish. This should get you to a reasonable level of purity.
One other thing that could be introducing error is your equipment. Are you using burettes, pippets and volumetric flasks in the proper manner to measure out all your liquids?



I am using all but a volumetric pippet.
I had been using a graduated cylinder to measure the acid, but I just received some 5mL volumetric pippets that I hope will help reduce the errors.

I have some Sodium carbonate recrystallizing as I write this.

One one more piece of laboratory equipment can help. That is a good lab scale. The one I use currently is accurate but I doubt it's precision.

Oscilllator - 2-3-2015 at 00:59

Quote: Originally posted by Atrum  

I am using all but a volumetric pippet.
I had been using a graduated cylinder to measure the acid, but I just received some 5mL volumetric pippets that I hope will help reduce the errors.

I have some Sodium carbonate recrystallizing as I write this.

One one more piece of laboratory equipment can help. That is a good lab scale. The one I use currently is accurate but I doubt it's precision.

It sounds like you're covering all your bases here;). One thing you could try to test if your scale is imprecise is to separately weigh two items, then weigh them together. If the separate weights of the two items do not add up precisely to the combined weight, then obviously something is off.
One other thing: How many times do you titrate each sample, and how much variance is there between each titre? I notice you don't mention averaging a number of trials.

Atrum - 3-3-2015 at 19:23

I perform at least 4. 1 to get a rough reading and 3 for the average.
The results are usually within 1 mL of each other. The problem is that is a pretty big difference in molarity when calculating.

AJKOER - 4-3-2015 at 11:41

Another idea when working with small quantities is to magnify the product. For example, react your H2SO4 with an excess of Zinc. Then capture and measure the volume of H2. Each mole of hydrogen formed implies a mole of Sulfuric acid consumed and will occupy 22.4 liters, or corresponding fraction thereof, per the reaction:

Zn + H2SO4 --) ZnSO4 + H2 (g)

This technique does require some knowledge of the starting strength of the acid (to access what constitutes an excess of Zinc) and assumes the reaction goes to completion (no sulfate coating inhibiting the reaction, for example). You can also dilute the acid employed in this experiment, to a limited degree, to lessen the percent error associated with visual reading of the volume of a concentrated acid although this may impact the speed of the hydrogen generation.

I would recommend working backwards from a targeted expected volume of hydrogen that can be conveniently collected on deciding the amount of acid to employ.
------------------------------

With respect to measurement error following a classical model, it has been shown that ( see, for example, https://www.google.com/url?sa=t&source=web&rct=j&... ):

β∗=β Var(X)/( Var(X)+Var(U))

where X, in our case, is a random variable relating to the concentration effect and U is the visual error (or measurement error) relating to reading volumes. Note, my suggestion on dilution by a factor of say k reduces the significance of the Var(U) by k square (as the Var(kX) = k squared * Var(X), and also the assumption that the visual reading error is not proportional to the entire volume of acid employed) and mitigates the bias usually present in measurement error situations involving the independent (or, so called predictor) variables. Also, averaging will not remove measurement based bias, but may reduce it (as if X is replicated 'n' times, we would now have the new term nVar(X)/(nVar(X)+Var(U)), which implies the significance of the Var(U) is effectively reduced by a factor of n).

One can also measure the value of Var(U)/Var(X) by several iterations of the experiment with an acid of known concentration.

[Edited on 5-3-2015 by AJKOER]

Atrum - 6-3-2015 at 07:11

Thanks AJKOER,

That sounds like an interesting way to determine concentration. May have to try it out.

blogfast25 - 6-3-2015 at 07:36

Quote: Originally posted by Atrum  
I perform at least 4. 1 to get a rough reading and 3 for the average.
The results are usually within 1 mL of each other. The problem is that is a pretty big difference in molarity when calculating.


Errors add up very quickly. You should get to a range of titrant volumes used (for say 3 titrations) of within 0.1 - 0.2 ml (for about 20 - 25 ml titrant used).

You need to:

1. use a decent Primary Standard
2. weigh to 1 mg or better
3. use Class A volumetric flasks for any sample or titrant preparation
4. use Class A volumetric pipettes
5. use a decent professional burette and make sure you read it properly (avoiding parallax errors)
6. familiarise yourself with end point determination
7. use a decent sized conical flask
8. make sure you titrate all of the sample: wash the walls of
your conical flask with small amounts of deionised water near the end point to make sure any splatter is accounted for
9. make sure you understand results calculations

Atrum - 6-3-2015 at 09:02

Quote: Originally posted by blogfast25  
Quote: Originally posted by Atrum  
I perform at least 4. 1 to get a rough reading and 3 for the average.
The results are usually within 1 mL of each other. The problem is that is a pretty big difference in molarity when calculating.


Errors add up very quickly. You should get to a range of titrant volumes used (for say 3 titrations) of within 0.1 - 0.2 ml (for about 20 - 25 ml titrant used).

You need to:

1. use a decent Primary Standard
2. weigh to 1 mg or better
3. use Class A volumetric flasks for any sample or titrant preparation
4. use Class A volumetric pipettes
5. use a decent professional burette and make sure you read it properly (avoiding parallax errors)
6. familiarise yourself with end point determination
7. use a decent sized conical flask
8. make sure you titrate all of the sample: wash the walls of
your conical flask with small amounts of deionised water near the end point to make sure any splatter is accounted for
9. make sure you understand results calculations


Hmm well my scale could be the big issue besides the purity of the primary standard I am using. It can only read to .01g.

Can anyone recommend a good scale that has a readability of .001g that doesn't cost $500 USD?

Atrum - 6-3-2015 at 09:06

Quote: Originally posted by blogfast25  
Quote: Originally posted by Atrum  
I perform at least 4. 1 to get a rough reading and 3 for the average.
The results are usually within 1 mL of each other. The problem is that is a pretty big difference in molarity when calculating.


Errors add up very quickly. You should get to a range of titrant volumes used (for say 3 titrations) of within 0.1 - 0.2 ml (for about 20 - 25 ml titrant used).

You need to:

1. use a decent Primary Standard
2. weigh to 1 mg or better
3. use Class A volumetric flasks for any sample or titrant preparation
4. use Class A volumetric pipettes
5. use a decent professional burette and make sure you read it properly (avoiding parallax errors)
6. familiarise yourself with end point determination
7. use a decent sized conical flask
8. make sure you titrate all of the sample: wash the walls of
your conical flask with small amounts of deionised water near the end point to make sure any splatter is accounted for
9. make sure you understand results calculations


Hmm well my scale could be the big issue besides the purity of the primary standard I am using. It can only read to .01g.

Can anyone recommend a good scale that has a readability of .001g that doesn't cost $500 USD?