The Lab of a HouseCat

Some chemistry experiments and observations even a house cat can do.

[housecatchem]

Burning magnesium

Metals can burn. Previously I’ve set steel on fire. Now it’s time to burn magnesium. Once ignited this metal gives a very bright flame and reaches very high temperatures. It is also hard to extinguish as it reacts not only with oxygen, but also nitrogen and carbon dioxide. 

[housecatchem]

Metal contest 3: Mn vs Zn

Here is a practical example of using the reactivity series. Which metal is more active? Manganese? Or zinc? Which one would replace the other from its solution?

For this experiment I used two test tubes. Zinc chloride solution is added to the first one and manganese chloride to the second one. A piece of manganese is added to the zinc chloride solution and a piece of zinc to the manganese chloride.

As you can see from the reactivity series, manganese comes before zinc, so it should replace zinc in zinc chloride and not vice versa:

K→Ca→Na→Mg→Al→Mn→Zn→Fe→Co→Ni→Sn→Pb→H→Cu→Ag

And indeed, there is nothing happening in the second test tube. But watch the first one...

There is a metallic zinc sediment forming around the piece of magnesium. Perhaps less spectacular than in case of tin or copper. Only visible when magnified. 

Mn (s) + ZnCl2 (aq) → Zn (s) + MnCl2 (aq)

You can also notice some gas bubbles. This is hydrogen formation as a result of a secondary process. Due to hydrolysis of zinc chloride in an aquatic solution, there is an acid environment containing hydrochloric acid. Manganese reacts with the acid which leads to the hydrogen gas release.

Hydrolysis (1st step):  Zn2+ + 2Cl- + HOH ⇄ ZnOH+ + Cl- + H+

Manganese reaction:  Mn (s) + HCl (aq) → H2 (g) + MnCl2 (aq)

[housecatchem]

Metal contest 2: copper dendrites

This experiment is another example of a single displacement reaction where a more active metal displaces the less active one from its solution. This time performed on a filter paper to reveal the beautiful copper crystals in a shape of a tree and thus called dendrites. 

For this experiment you’ll need a 10% solution of a copper sulfate, a filter paper in a dish (like a Petri dish) and some pieces of metals. A zinc pellet or a piece of iron is good (both metals are more active than copper). It’s also possible to use different metals and compare the result. There intensity of dendrite growth will depend on the metal activity.

If you have a MEL Science subscription, you might already have a dish and a zinc pellet from the “Tin” set and a copper sulfate from the “Metal contest” set. To make a solution, dissolve a half of a big measuring spoon of copper sulfate in 10 ml water. As for filter paper, you can take any paper that absorbs water. A coffee filter or even a paper kitchen towel. As for the pieces of metals, try a paperclip, a nail, a pin or even a hedgehog from the “Tin hedgehog” experiment. Just take into account, that paperclips may be made of different metals, not only iron. And steel (mostly iron) nails, pins and paperclips may be covered with a thin layer of zinc. It is not a problem for this experiment. Make sure that the metallic object has a flat side to make good contact with a filter paper.

The instructions are easy. Place the filter paper into a dish and moisten it with the 10% copper sulfate solution. Not too much, just make it all wet. Place the pieces of metal on the filter paper. Make sure they touch the paper. Cover the dish and wait for some hours (or leave it for the night). Observe the brownish dendrites of metallic coppers grown around the piece of metal.

In case of several pieces of different metals, the result will reflect the metal’s reactivity to some extent.

In general, the more active metal (see the reactivity series), the more dendrites should grow and the farther they reach. But even there is a correlation between the metal’s reactivity and the amount of dendrites, there are still other factors to consider. Like the size of the pellet, its purity, its ability to form a protective layer of oxides (which prevents the reaction), contact between the metal surface and the solution on the filter paper etc.

You may also like to watch the MEL Science version of this experiment (see instructions in the video description on YouTube):

[housecatchem]

Metal contest

MEL Science suggests to look at the single displacement reaction as a sport competition where the metals are athletes. In this contest the more active metal replaces the less active one in a solution. And the “strength” or reactivity is pre-determined and arranged in a reactivity series.

In their experiment from the “Metal contest” set they suggest to observe how more active zinc replaces copper from copper (II) sulfate solution and tin from tin (II) chloride. You can find the detailed instructions and explanation on their webpage.

An easy way to observe a single displacement reaction. There is not much to add, so here is what I got:

The reactivity series gives a lot of combinations to experiment with, so there is more to come. And  at risk of preempting myself, let me tell you that not all of them will go as expected.

Like on this video you can clearly observe gas bubbles as a result of a secondary process. Both copper (II) sulfate and tin (II) chloride tend to hydrolyse in a solution and form acidic environment (sulfuric or hydrochloric acid):

Cu2+ + SO42- + HOH ⇄ CuOH+ + SO42- + H+ 

Sn2+ + 2Cl- + HOH ⇄ SnOH+ + Cl- + H+ 

Zinc reacts with either sulfuric or hydrochloric acid and replaces hydrogen which forms gas bubbles and escapes.

[housecatchem]

What’s in the tin?

Experimenting with tin made me curious to the properties of this element. So I decided to put together a collection of videos on tin’s properties.

From Periodic Videos:

From Thoisoi:

From MEL Science:

There is a text version for this video with some more details: https://melscience.com/NL-en/articles/tin-crackling-metal/

and an extra bonus from them: https://melscience.com/NL-en/articles/tin-its-oxidation-states-and-reactions-it/

[housecatchem]

The hedgehog reversed: dissolving tin

What can you do with the tin hedgehog from one of the previous experiments? One of the possibility is to study the properties of tin. Like its interaction with acids and alkali. And at the same time to get tin salts for further experiments.

The following experiments with acids and alkali are less safe than those from the MEL Science sets (if you use any) and require some extra chemicals.

To “reverse” the tin hedgehog creation and turn it back to the tin chloride, it’s possible to dissolve it in hydrochloric acid. As the reaction is slow, it’s better to use higher concentrations of the acid (I used 30%) and heat (putting the test tube into the glass of hot water is sufficient). You can find the hydrochloric acid (also called muriatic acid) in a hardware store. Mostly it’s about 10%, so it will take more time for tin to get dissolved.

What are the other ways to dissolve the tin hedgehog? Tin can interact with both acids and alkali.

This video shows the tin interaction (or the absence of it) with hydrochloric acid, sulphuric acid and potassium hydroxide.

In the reactivity series tin comes before the hydrogen, so it should react with acids to produce hydrogen gas and tin salts. But at the same time tin is placed is in the middle of the series, close to the hydrogen. It is less reactive metal than for example magnesium or zinc, so its reaction with acids will go much slower.

K→Ca→Na→Mg→Al→Zn→Fe→Co→Ni→Sn→Pb→H→Cu→Ag

Hydrochloric acid is the best way to dissolve tin as you can see from both videos. Tin reacts with the diluted acid, but very slowly. Concentrated acid dissolves tin much quicker, especially when heated. The product of this reaction tin chloride interacts with hydrochloric acid and forms complex compounds.

Sn (s) + 2HCl (aq) → SnCl2 (aq) + H2 (g)

SnCl2 (aq) + HCl (aq) → H[SnCl3] (aq)

The reaction with a diluted sulphuric acid is very slow even at higher temperatures - I could hardly notice any hydrogen bubbles. The concentrated sulphuric acid slowly dissolves tin when heated (oxygen from the air seems to help too):

Sn (s) + 4H2SO4 (conc)  → Sn(SO4)2 (aq) + 2SO2 (g) + 4H2O

Sulphuric acid can be bought at 37% concentration (used as a battery acid), but due to the lack of interaction, you can’t use it to dissolve tin.

Interaction with nitric acid varies depending on the acid concentration and the temperature. 

Interaction with a very diluted nitric acid (3-5%) at low temperatures leads to tin nitrate formation without any visible gas release:

4Sn (s) + 10HNO3 (very diluted)  → 4Sn(NO3)2(aq) + NH4NO3 (aq) + 3H2O

Higher concentrations of nitric acid and higher temperatures lead to nitrogen formation (with some impurities of nitrogen monoxide):

5Sn (s) + 12HNO3 (diluted)  → 5Sn(NO3)2(aq) + N2 (g) + 6H2O

Concentrated nitric acid oxidizes tin to a stannic acid:

Sn (s) + 4HNO3 (conc)  → SnO2•xH2O (s) + 4NO2 (g) + (2-x)H2O

Unfortunately nitric acid is not readily available and its use is restricted, so no demonstration in a house lab. But there is a video by Tuan Dang:

Tin can be also dissolved in sodium or potassium hydroxide. The result may vary depending on temperature and concentration:

Sn (s) + KOH (conc) + 2H2O  → K[Sn(OH)3] (aq) + H2 (g)

Sn (s) + 2KOH (conc) + 4H2O  → K2[Sn(OH)6](aq) + 2H2 (g)

[housecatchem]

Foil etching experiment

The “Metal contest” set by MEL Science contains a “Foil etching” experiment. An interesting reaction which appears to be not easy to explain. A mix of copper sulphate and sodium chloride in water starts to “eat through” the layer of aluminium foil as if it was a “dangerous acid“ (a nice experiment for a “Hollywood effect“):

The explanation seems to be simple - it is an example of a single displacement reaction. Aluminium is more reactive than copper, so aluminium metal can replace copper in its solutions. But why is sodium chloride added? And can the reaction go without it? Let’s try it out:

It looks like chloride ions are essential here. Their effect can be explained as followed. Aluminium reacts with oxygen from the air and gets covered with a thin film of aluminium oxide. This film protects the metal from further oxidation and other interactions (with water or salts). Protected by this film aluminium foil doesn’t react with a copper sulphate solution. Some sources mention that the reaction is taking place, but very slowly (the foil would eventually dissolve in a couple of days).

Chloride ions somehow compromise this protection and without its shield aluminium foil starts to react with copper sulphate. But why do we observe the effervescent hydrogen? Well, without its protective film aluminium is a very reactive metal and can interact with water in a similar way as alkali or earth alkali metals do. 

So there are two main reactions are taking place here:

2Al + 6H₂O → 2Al(OH)₃ + 3H₂

3Cu²⁺ + 2Al → 3Cu + 2Al³⁺

It’s interesting to notice that aluminium doesn’t react with a sodium chloride solution. Somehow both chlorides and copper are needed for this process to go. More investigation is needed to learn the nature of this reaction.

[housecatchem]

Growing tin dendrites

A single displacement reaction (as in case of tin hedgehog or silver fur) is not the only method to throw out a metal of its solution. Electrolysis is another one. When current flows through a solution, the electrons connect to the ions and reduce them to a metal.

In case of tin (II) chloride you can see the metallic tin disposing on a cathode. Tin ions are reduced to the metallic tin and form dendrite like crystals. The metal of anode is loosing its electrons, turns into ions and “dissolves”. The growth of tin dendrites is a beautiful process to observe.

For this experiment I followed the instruction from the MEL science webpage and used the supplies from the “Tin” set. Please, visit their page for more information and a detailed instruction video.

As you can see, I have performed this experiment twice (the first video was not sharp enough as you may have noticed). For the second time I used paperclips as electrodes to prevent the dissolving of a crocodile clip.

The chemistry of this process is already described in details on the MEL Science page shown above. Shortly:

Cathode-: 2Sn2+(solution) + 4e-→ 2Sn(solid)

Anode+: 2H2O – 4e- → O2 + 4H+

And a secondary reaction:  2SnCl2 + O2 + 2H2O → 2SnO2 + 4HCl

[housecatchem]

More of silver crystals

A collection of videos I found on Internet on a single displacement reaction involving silver.

1. Chemistry in Context has made an opera about this reaction:

2. A silver Christmas tree from the MEL Science:

And some instructions: https://melscience.com/en/articles/silver-tree-experiment/

3. A time lapse video from Jarno Photography:

4. A video from the NurdRage exploring a different way of growing silver crystals by means of electrolysis: 

[housecatchem]

Silver Fur Snake

After growing a tin hedgehog let’s add another “exotic animal” to our chemical zoo. A silver fur snake for example. Made from a copper wire and silver nitrate solution.

It’s the same type of interaction, a singe displacement reaction based on the metal activity. Copper as a more active metal can replace silver in its salts. Metallic silver starts to grow on the copper wire as a grey fur. But there are also shiny metal crystals to see.

For this experiment I added a 1,5% solution of silver nitrate into a test tube with a copper wire. Lower concentrations of silver nitrate may result in a slower growth, but better qualities of the crystals.

Chemistry behind this process:

[housecatchem]

Reducing properties of Tin (II)

While growing the tin hedgehog from the previous experiment, let’s learn some other properties of tin (II) chloride. It is a strong reducing agent. Tin (II) ion interacts with the oxidizers like bromine water, potassium permanganate, potassium bromide. And also reduces iron (III) ion to a bivalent iron or silver (I) ione to metallic silver.

For this experiment I took iodine solution in alcohol (instead of bromine water), potassium permanganate, potassium bichromate and iron (III) chloride (with KSCN added for a better visual effect, you can find more information here).

Those solutions were dropped on a filter paper resulting in four colored spots. Tin (II) chloride solution was dropped in the middle of each spot which resulted in discoloration as you can see from the video:

In case of bichromate reduction, a light green color of chrome (III) can be observed. Manganese (II) salts have a light pink color in a solid state, but it’s hard to notice in a solution. On a side note - try to compare a manganese (II) salt solution with water under a good light and you may see the difference.

The chemistry behind the discoloration can be generally described as followed:

[housecatchem]

Tin crystals as a hedgehog

This is an example of a displacement reaction. A more active element displaces a less active one from its salts. In case of metals there is a reactivity series where all metals are arranged by their activity from the highest to the lowest (vertically or horizontally).

A short example of reactivity series looks like:

K→Ca→Na→Mg→Al→Zn→Fe→Co→Ni→Sn→Pb→H→Cu→Ag

Potassium is the most active in this particular example and silver is the least one.

For a full version, please, visit a Wikipedia article on reactivity series.

In this experiment zinc as a more active metal displaces tin from a solution of tin chloride. Metallic tin is disposed on a zinc pellet and forms beautiful crystals. A tin covered zinc pallet might remind you of a hedgehog. Or a lovely frosted pine tree (with help of a macro lens).

The reaction can be written as:

SnCl2 + Zn → Sn + ZnCl2 

Addiction of acid to tin chloride solution is needed to suppress its hydrolysis in water. Normally a hydrochloric acid is used for a chloride salt, but other strong acids will do too. Even an acidic salt like sodium hydrogen sulphate (which is safer to work with than an acid).

I performed this experiment with the help of MEL chemistry “Tin” set. In their sets they provide all materials needed to perform 2 or 3 experiments.

On their page you can find more detailed explanation of this experiment: 

https://melscience.com/en/experiments/tin-hedgehog/

And this is how my “tin hedgehog” looks like after 15 minutes:

[housecatchem]

The colors of copper vitriol

Another illustration to the experiment of yesterday: 

In descending order:

Copper sulphate pentahydrate.

Copper sulphate anhydrous.

Copper sulphate anhydrous after being exposed to ammonia gas.

Adding water destroys the ammonia complex. The basic copper sulphate is formed (among others).

[housecatchem]

Copper sulphate dehydration

A blue pentahydrate is the most common form of copper sulphate. When heated it looses water and becomes anhydrous (you can observe it turning into a grayish powder). Adding water turns it back to the blue hydrated form. This is probably one of the easiest and the most performed experiment. The reaction can be described as followed:

CuSO4·5H2O  →   CuSO4 + 5H2O

But anhydrous copper sulphate can also combine up to 5 ammonia molecules when introduced to a dry ammonia gas. In the presence of water both ammonia and water can be absorbed.

With dry ammonia the violet blue CuSO4·5NH3 is formed.

In the presence of water (moist):  CuSO4·4NH3·H2O and CuSO4·2NH3·3H2O

Within the house lab I could not work completely water free, so I decided to cover the anhydrous copper sulphate with a glass (to prevent it from contacting water of the ammonia solution), add a couple of drops of 25% ammonia onto the glass and cover the dish (to keep ammonia inside). On the video you can see how the color changes into violet blue as ammonia gas reaches the powder. Unfortunately it takes too long. A better alternative is to take two cotton pads. Put a dry one on the powder (to protect it from water), drop a little bit of a strong ammonia solution on a second one, place it above the first one and cover the dish with a lid. It takes about 5-10 minutes for the powder to turn violet blue.

Adding water destroys the ammonia complex and forms the basic form of copper sulphate (you notice the turquoise substance is not as soluble in water as copper sulphate). 

More details:

Dehydration of copper sulphate takes several steps:

At around 93°C (other sources mention 63°C) two water molecules split off forming  CuSO4·3H2O

At 105°C (109°C is also mentioned), the monohydrate CuSO4·H2O is formed.

Heating up to 230°C (numbers like 200°C and 220-250°C are also mentioned) leads to the anhydrous form.

Above 650-700°C CuSO4 decomposes:

2CuSO4 → 2CuO + 2SO2 + O2  (probably via formation of  Cu2OSO4 ).

Copper sulphate pentahydrate naturally occurs as Chalcanthite:

From: Planet Mine, La Paz County, Arizona, United States of America

Picture Source: e-rocks.com.

A heptahydrate CuSO4·7H2O is known as a rare mineral Boothite:

From: Copper Basin, Battle Mountain District, Lander Co., Nevada, USA 

Picture Source: mindat.org.

CuSO4·3H2O occurs as Bonattite:

From: La Compania Mine,Chile

Picture source: mineralatlas.eu.

Anhydrous CuSO4 occurs as Chalcocyanite:

From: Ronneburg U deposit, Gera, Thuringia, Germany

Picture Source: mindat.org.

[housecatchem]

Alkali metals in water

All schoolbooks mention the reaction of alkali metals with water. Sometimes even a demonstration is given during the chemistry class. But these experiments can be quite dangerous to perform at home. Fortunately there is YouTube:

Here is an older movie of a lower quality showing the reaction of all alkali metals except Francium: 

There is also a better quality video that even shows how would Francium react with water:

[housecatchem]

Native fluorine

Fluorine is a very reactive element. It forms compounds with almost all other elements. Only some noble gases are the rare exception. That is why fluorine is impossible to find in its free form, right?

But in 2012 the free fluorine was finally found in nature. And at the same time the mystery behind the “stinky rocks” (also known as Stinspat or fetid fluorite) was solved. Meet Antozonite, a variety of fluorite.

A picture from e-rocks.com shows Antozonite (a miniature  6.2 × 6.7 × 2.6 cm) found at Wölsendorf, Bavaria, Germany.

Antozonite seems to have multiple inclusions filled with pure fluorine, which possess an unpleasant ozone of chlorine like smell. Once crashed, antozonite frees fluorine which immediately reacts with oxygen and water vapors from the air and forms ozone and hydrogen fluoride. All together those compounds give the mineral an unpleasant smell after crashing. But for a long time people believed that it was a mythical substance “antozone” which was responsible for the smell. 

Antozonite is a radioactive form of fluorite (calcium fluoride) containing some uranium inclusions. Uranium is responsible for emitting beta radiation which turns the stone dark purple and breaks some of calcium fluoride apart. The fluorine feels the pockets inside the stone where it can be stored as it doesn’t react with fluorite.

Another interesting picture of this mineral from e-rocks.com:

Here you can see “Uraninite crystals and antozonite with weathered hornblende. An interesting suite of representative materials from this locality.” Found at Tripp Nu-Age Mine, Cardiff Township, Ontario, Canada.

More on this mineral:

1. Wikipedia 2. mindat.org 3. An article explaining the origin of smell in Atnozonite (in Nature.com)

[bijoux-et-mineraux]

Rare Phosphophyllite - Unificada Mine, Potosi, Bolivia

[housecatchem]

A rare mineral Phosphophylite is a hydrated zinc phosphate, but also contains iron (II) and manganese with a formula  Zn2Fe(PO4)2•4H2O or  Zn2Fe2+0.75Mn2+0.25(PO4)2•4(H2O). It’s often associated with vivianite, phosphosiderite and triphylite among others (chalcopyrite, sphalerite). 

More information:

Wikipedia

Mineralienatlas

Webmineral

Mindat.org