#houselab

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+

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):

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)

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