{"id":55,"date":"2013-04-24T01:14:46","date_gmt":"2013-04-23T15:14:46","guid":{"rendered":"http:\/\/www.trentwallis.com\/chemblog\/?p=55"},"modified":"2017-08-04T17:07:43","modified_gmt":"2017-08-04T07:07:43","slug":"titrations-at-home","status":"publish","type":"post","link":"http:\/\/www.trentwallis.com\/chemblog\/titrations-at-home\/","title":{"rendered":"Titrations at home"},"content":{"rendered":"

This is a #realtimechemcarnival<\/a> post for #realtimechem<\/a>\u00a0week. If you’re not familiar with #realtimechem, you can read the FAQ here<\/a>.<\/p>\n

I wanted to present an example of a titration that can be performed at home, using chemicals that you might have in your kitchen\/pantry\/wherever you keep your cleaning stuff. In my case, I wanted to determine the purity of sodium hydroxide by titration against citric acid.<\/p>\n

What, you don’t have any volumetric glassware at home? Well me neither. \ud83d\ude42<\/p>\n

Equipment<\/strong><\/h2>\n

In the absence of any volumetric glassware, it is possible to measure solutions by mass (using an electronic balance) rather than by volume. My balance\u00a0only reads to the nearest 0.1 g, so this titration is not going to\u00a0be as accurate as a skilled operator in the lab, but it’s not bad. If your balance only reads to the nearest gram then you could still use this method but your accuracy will be limited by your balance unless you use very large masses.\u00a0By the way, you can get electronic balances that read to 0.01 g on Ebay for under $10 shipped; these would ensure that your accuracy is not limited by your balance.<\/p>\n

The only other equipment I used were some plastic drinking cups and plastic straws.<\/p>\n

Chemicals<\/strong><\/h2>\n

I weighed out 16.0 g of powdered NaOH (which I have used for unclogging drains and usually resides under my sink). It’s a couple of years old.<\/p>\n

\"TWP_4064\"<\/a><\/p>\n

I then dissolved the 16.0 g NaOH in tap water to give a solution with a total mass of 361.1 g. In molar terms, this is a bit over 1 M. If I had some deionised water around I would have used that, but I didn’t.<\/p>\n

\"TWP_4065\"<\/a><\/p>\n

A standard solution of citric acid was then prepared in a similar way: 25.9 g of citric acid in 384.8 g solution.<\/p>\n

\"TWP_4066\"<\/a><\/p>\n

Indicator<\/strong><\/h2>\n

The titration curve for my citric acid-sodium hydroxide titration should look something like this (I generated this with the wonderful\u00a0CurTiPot<\/a>\u00a0using my approximate reagent concentrations and volumes):<\/p>\n

\"titration<\/a><\/p>\n

As an aside, the acidic region of this graph is somewhat unusual: the (almost) linear dependence of pH on volume in this region is a result of citric acid being a weak triprotic acid whose pKa values are quite close together.<\/p>\n

Now I don’t have a pH or conductivity meter at home and although it is possible to build<\/a>\u00a0your own<\/a>\u00a0I’m going to go with a halochromic<\/a> indicator. In the lab phenolphthalein<\/a> would usually be the indicator of choice, but any indicator that changes colour between pH 8 and 12 will do.<\/p>\n

It’s very easy to make an indicator solution at home, indeed a common school practical is to make an indicator solution by extracting compounds from black tea or red cabbage. I was using some\u00a0purple carrots<\/a>\u00a0for dinner, so I finely chopped the tops and tails and boiled them with water in the microwave to extract the anthocyanins<\/a>, leaving a deep purple<\/a> solution.<\/p>\n

I used the tops and tails of about 8 purple carrots, but you could get away with using far less than this. I obtained about 300 mL of indicator solution using multiple extractions (total overkill) and I could easily have extracted much more.<\/p>\n

\"TWP_4061\"<\/a><\/p>\n

Titration<\/strong><\/h2>\n

The titration is performed in the usual way, adding one reagent to a measured sample of the other. The only difference in this procedure is that the solutions are measured by mass, rather than by volume. I transferred all solutions to the reaction vessel<\/del>\u00a0plastic cup using plastic drinking straws with my index finger over the top.<\/p>\n

I started with about 15 g (measured exactly, see results below) of the citric acid solution, then added a couple of mL of indicator, which turns red.<\/p>\n

\"TWP_4071\"<\/a><\/p>\n

The drops on the side of the flask are water from washing the cup after the previous titration.<\/p>\n

I then zeroed the balance, removed the cup from the balance and started to add the NaOH solution. After about 10 g of the NaOH solution was added, the solution had turned a pink colour:<\/p>\n

\"TWP_4075\"<\/a><\/p>\n

The solution turned a purple colour near the endpoint:<\/p>\n

\"TWP_4078\"<\/a><\/p>\n

These colour changes (from red to pink to purple) are gradual, as expected from the gradient of the graph above.<\/p>\n

Endpoint occurs when one or two drops of the NaOH solution turn the reaction mixture from purple to a blue-black colour:<\/p>\n

\"TWP_4084\"<\/a><\/p>\n

Results<\/strong><\/h2>\n\n\n\n\n\n\n\n\n\n\n
\n

m(CA sol’n)<\/p>\n<\/td>\n

\n

m(NaOH sol’n)<\/p>\n<\/td>\n

\n

NaOH:CA mass ratio<\/p>\n<\/td>\n<\/tr>\n

\n

15.7<\/p>\n<\/td>\n

\n

17.2<\/p>\n<\/td>\n

\n

1.10<\/p>\n<\/td>\n<\/tr>\n

\n

16.0<\/p>\n<\/td>\n

\n

17.6<\/p>\n<\/td>\n

\n

1.10<\/p>\n<\/td>\n<\/tr>\n

\n

15.7<\/p>\n<\/td>\n

\n

17.3<\/p>\n<\/td>\n

\n

1.10<\/p>\n<\/td>\n<\/tr>\n

\n

15.7<\/p>\n<\/td>\n

\n

17.1<\/p>\n<\/td>\n

\n

1.09<\/p>\n<\/td>\n<\/tr>\n

\n

15.5<\/p>\n<\/td>\n

\n

17.2<\/p>\n<\/td>\n

\n

1.11<\/p>\n<\/td>\n<\/tr>\n

\n

17.2<\/p>\n<\/td>\n

\n

19.0<\/p>\n<\/td>\n

\n

1.10<\/p>\n<\/td>\n<\/tr>\n

\n

15.9<\/p>\n<\/td>\n

\n

17.6<\/p>\n<\/td>\n

\n

1.11<\/p>\n<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n

The method seems quite reproducible within the limitations imposed by the equipment.<\/p>\n

Calculations<\/strong><\/h2>\n

[citric acid] =\u00a025.9 g \/ 384.8 g = 67.3 mg of citric acid per gram of solution<\/p>\n

Converting from mg to mmol:<\/p>\n

[citric acid] = 67.3 \/ 192.124 = 0.350 mmol of citric acid per gram of solution<\/p>\n

Average ratio of m(NaOH solution):m(citric acid solution) = 1.10<\/p>\n

[NaOH] = 0.350 * 3 \/ 1.10 = 0.955 mmol of NaOH per gram of solution<\/p>\n

m(NaOH) = 0.955 * 39.998 = 38.2 mg of NaOH per gram of solution<\/p>\n

m(NaOH total) = 38.2 * 361.1 g = 13.8 g NaOH<\/p>\n

% purity = 13.8 \/ 16.0 = 86.2%<\/p>\n

Discussion<\/strong><\/h2>\n

Titrating by mass rather than by volume is straightforward, fast and reproducible, even with the limitations of kitchen\/household apparatus. I am quite keen to perform this with more precise drop control, with a balance reading to 0.01 g, and with phenolphthalein as an indicator.<\/p>\n

Pedagogically, it is probably not as instructive as volumetric analysis but it is certainly an interesting variation.<\/p>\n

One thing I did wonder was if my “citric acid” was anhydrous (as assumed) or the monohydrate, but no change in mass occurred after heating 10.6 g of it in the oven at 130 \u00b0C for about an hour, which suggests that it is indeed anhydrous.<\/p>\n

Any comments are greatly appreciated!<\/p>\n

\"Creative<\/a>
\nThis work is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 3.0 Australia License<\/a><\/p>\n","protected":false},"excerpt":{"rendered":"

This is a #realtimechemcarnival post for #realtimechem\u00a0week. If you’re not familiar with #realtimechem, you can read the FAQ here. I wanted to present an example of a titration that can be performed at home, using chemicals that you might have … Continue reading →<\/span><\/a><\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[2,3],"tags":[],"class_list":["post-55","post","type-post","status-publish","format-standard","hentry","category-chemistry","category-practical"],"_links":{"self":[{"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/posts\/55","targetHints":{"allow":["GET"]}}],"collection":[{"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/comments?post=55"}],"version-history":[{"count":10,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/posts\/55\/revisions"}],"predecessor-version":[{"id":486,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/posts\/55\/revisions\/486"}],"wp:attachment":[{"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/media?parent=55"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/categories?post=55"},{"taxonomy":"post_tag","embeddable":true,"href":"http:\/\/www.trentwallis.com\/chemblog\/wp-json\/wp\/v2\/tags?post=55"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}