copper carbonate and hydrochloric acid experiment

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Copper and Nitric Acid Chemistry Demonstration

The copper and nitric acid reaction is a dramatic color change chemistry demonstration. The reaction illustrates several chemistry principles, including exothermic reactions , redox reactions , coordination complexes, oxidation, oxidation states , and the metal activity series . Here are instructions explaining how you perform this demonstration safely, with a look at its chemical reactions.

You only need two common chemicals. The most important part of the reaction is the choice of reaction vessel. The reaction produces heat, so use a study glass container.

• 40 ml concentrated nitric acid (HNO 3 )
• 1-liter flask (Erlenmeyer, boiling flask, or Buchner flask)
• Clamp stand
• Bowl (optional)

The original demonstration uses a copper penny, but modern pennies are zinc plated with a thin layer of copper. A better choice is a piece of copper wool or some copper shavings. The reaction works fine with copper wire, but is not as dramatic because the wire has less surface area.

A smaller version of the demonstration uses a bit of copper, a small volume of nitric acid, and a borosilicate glass test tube.

Perform the Copper and Nitric Acid Chemistry Demonstration

Nothing could be easier! Set up and perform the demonstration inside a fume hood.

• Pour the nitric acid into the flask.
• When you are ready for the reaction, add the copper.

Initially, the nitric acid attacks the copper, turning the liquid green and releasing heat and reddish brown nitrogen dioxide vapor. Eventually, even the liquid turns brown.

• Add water and dilute the solution.

Diluting the acid changes the conditions. The liquid changes color into a bright blue, while the vapor changes from reddish brown to colorless.

A Look at the Chemistry

If you look at the metal reactivity series, copper is pretty unreactive. It’s even considered a noble metal by some chemists. It resists oxidation by hydrochloric acid (HCl), yet readily reacts with nitric acid (HNO 3 ). This is because nitric acid acts both as an oxidizer and an acid. Copper reacts with nitric acid, forming aqueous copper nitrate, nitrogen dioxide gas, and water.

Cu(s) + 4HNO 3 (aq) → Cu(NO 3 ) 2 (aq) + 2NO 2 (g) + 2H 2 O(l)

The reaction immediately produces heat (reaching 60 to 70 degrees C) and releases deeply-colored nitrogen dioxide gas. The green color comes from copper(II) ions forming a coordination complex with nitrate ions. Diluting the concentrated acid with water changes the liquid color to blue as the water displaces the nitrate ions, leaving only aqueous copper(II) nitrate. The water reacts with nitrogen dioxide and forms nitric oxide.

3Cu(s) + 8HNO 3 (aq) → 3Cu 2+ (aq) + 2NO(g) + 4H 2 O(l)+ 6NO 3 − (aq)

The concentration of the acid affects its oxidizing capacity. For example, copper does not react with dilute sulfuric acid (H 2 SO 4 ), but a similar reaction occurs in concentrated sulfuric acid:

Cu + 2H 2 SO 4  → SO 2  + 2H 2 O + SO 4 2−  + Cu 2+

Containing the Copper and Nitric Acid Reaction

A few simple revisions contain the reaction and improve both the safety and dramatic effect of the copper and nitric acid chemistry demonstration. You can perform this variation of the copper and nitric acid reaction out in the open, but it’s still a good idea to separate the set-up from the audience using a safety shield.

• Add nitric acid to a round-bottomed borosilicate flask. Clamp it into position on a stand. Ideally, use a borosilicate flask and place a bowl beneath the flask in case the glass leaks or breaks.
• Fill an Erlenmeyer (conical) flask with water and clamp it into position near the round flask.
• Stopper the round flask (acid) and loosely plug the conical flask with glass wool. The glass wool prevents the escape of nitrogen dioxide into the outside air. Insert glass tubing the ends reach the bottoms of each flask. (Don’t use plastic tubing.)
• When you are ready for the demonstration, add the copper to the borosilicate flask and fit the stopper and tube onto it.

Initially, the liquid in the round flask turns green and evolves reddish brown nitrogen dioxide. After about a minute and a half, the reaction slows and cools. The pressure reduction from the cooling draws water in from the conical flask. This dilutes the nitric acid and also reacts with the nitrogen dioxide gas, forming a fountain. Finally, the liquid in the round flask turns blue as copper nitrate forms.

Safety and Disposal

• Only perform this demonstration if you are a chemist or chemistry educator and have access to proper safety gear and a fume hood. Nitric acid is a corrosive strong acid, while nitrogen dioxide is a toxic reddish-brown gas. Wear gloves, goggles, and a lab coat. Perform the open demonstration under a fume hood.
• Please choose sturdy glassware for this demonstration. The initial reaction produces heat, so there is a risk of glassware breakage. For this reason, a boiling flask is ideal. Alternatively, use a Buchner flask.
• After the demonstration, neutralize the dilute nitric acid using any inorganic base, such a baking soda, sodium hydroxide solution, or potassium hydroxide solution. The neutralization reaction also produces some heat. Afterward, you can safely wash the liquids down the drain with water.
• Cotton, F. Albert; Wilkinson, Geoffrey (1988). Advanced Inorganic Chemistry (5th ed.). New York: John Wiley & Sons. 769-881.
• Shakhashiri, Bassam Z. (1985). “Properties of Nitrogen(II) Oxide”. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 2 . The University of Wisconsin Press. ISBN: 978-0299101305.
• Shakhashiri, Bassam Z. (1985). “Coin-Operated Red, White, and Blue Demonstration: Fountain Effect With Nitric Acid and Copper”. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 3 . The University of Wisconsin Press. 83-91. ISBN: 978-0299119508.
• Summerlin, Lee R.; Borgford, Christie L., Ealy, Julie B. (1988 ). Chemical Demonstrations: A Sourcebook for Teachers Volume 2 (2nd ed.). American Chemical Society. ISBN: 978-0841215351.

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12.4 The general reaction of an acid with a metal carbonate

Reactions of acids with bases

Chapter overview

The central challenge of this chapter is to establish that acid-base reactions are exchange reactions. A fragment of the acid is exchanged with a fragment of the base and a salt and water are the resulting products of the reaction. The type of salt that forms depends on the identities of the acid and the base that were combined during the reaction.

Once learners understand this, they have taken an important step to understanding acid-base chemistry. We will spend some time developing a frame for explaining this at the start of the chapter, to which we will return frequently.

In light of the fact that learners have yet to learn about cations and anions, we have considered it pedagogically justifiable to make the following simplifications to currently accepted acid-base theory, in order to bring the concept of exchange across to the learners:

Acids can be thought of as contributing H (instead of H+); and

Bases can be thought of as contributing O or OH (instead of O 2 - and OH-).

Water (H 2 O) is a combination of 2 H and 1 O, or alternatively 1 H and 1 OH.

We are well aware that writing H + OH → H 2 O has no meaning in science and for this reason we have avoided this usage in the text. But we do consider the use of simplified symbols (H instead of H+ and so forth) to have an advantage over their scientifically correct (but potentially confusing) counterparts in this context.

There is also a danger that misconceptions and sloppy usage of symbols may result further down the line, when simplifying in this way. However, we feel these risks are counterbalanced by the greater likelihood of learners understanding the concept of exchange if the symbols they work with are not cluttered with additional information - like the charges on the ions - that have no meaning for them yet.

Other skills that will be reinforced in this chapter are:

• writing chemical formulae;
• converting between word equations and chemical equations; and
• balancing chemical equations.

A word of caution: Acid-base reactions are neutralisation reactions. However, this does not mean that the mixture of an acid with a base will be a neutral solution and you should avoid language that reinforces this notion. Even if equivalent quantities (stoichiometric quantities) of the acid and base are mixed - which would imply that both have been neutralised - the resulting solution will only be neutral (i.e. pH = 7) under very special circumstances. The reason is that not all salts are 'neutral substances'; in fact most salts have acid-base properties of their own. The chemistry required for learners to understand this is beyond them at this stage and will only be dealt with in Physical Sciences in Grade 12. Our suggestion is that you simply refrain from calling salts 'neutral substances'. If questions arise around the issue you could point out that the salts they will encounter in this chapter may be neutral substances, but that this is not true of all salts.

Take note that although there is no section specifically named 'Applications' as indicated in CAPS, this content has rather been dealt with under other sections where it is more appropriate.

6.1 Neutralisation and pH (1.5 hours)

6.2 The general reaction of an acid with a metal oxide (1.5 hours)

6.3 The general reaction of an acid with a metal hydroxide (1.5 hours)

6.4 The general reaction of an acid with a metal carbonate (1.5 hours)

• What is the reaction between an acid and a base called?
• What happens to the pH when an acid and a base are mixed?
• Does the reaction between an acid and a base always give a neutral mixture, in other words a mixture with pH = 7?
• Which factors will determine the pH of the final solution when an acid and a base are mixed?
• Is there a way to predict which classes of compounds will tend to be acids and which will tend to be bases?
• Are metal oxides, metal hydroxides and metal carbonates acidic or basic? Which pH range will their solutions fall into?
• What products can we expect when a metal oxide, a metal hydroxide or a metal carbonate react with an acid?
• Are there general equations to explain these reactions?
• How does acid rain form?

Neutralisation and pH

• neutralisation reaction
• neutral solution
• exchange reaction
• laboratory acids

In the previous chapter we learnt about a new concept, namely pH. If we want to know whether something is an acid or a base, we can measure its pH:

• Acids have pH values below 7. The lower the pH value, the more strongly acidic the substance.
• Bases have pH values above 7. The higher the pH value, the more strongly basic the substance.
• Neutral substances have pH equal to 7.

Another useful thing we learnt in the previous chapter is that we can use universal indicator to measure the pH of a solution. Universal indicator has different colours at different pH values. Below is a colour chart showing the range of colours for universal indicator and the pH values they correspond to. You will need it for all the activities of this chapter, because we are going to do lots of pH measurements!

Can you remember how we used the universal indicator paper in the previous chapter? Here are some suggestions for the investigations of this chapter:

• Before you start, place 1-cm lengths of universal indicator paper on a sheet of white paper, like this:

Later, if you want to write down a note or an observation, you can do so directly on the paper and copy it to your workbook afterwards.

Instead of dipping the paper in the solutions you are testing, use a glass rod or drinking straw to transfer a drop of the test solution to the indicator paper.

For some of the investigations in this chapter, you will be using droppers or syringes to measure out quantities. Tell learners that they may not use droppers or syringes to squirt water at other learners! There are many reasons why this is not a good idea. The most important reason is that the dropper or syringe may contain acid , that could end up in someone's eye where it could cause permanent damage or even blindness. So, squirting each other with the droppers or syringes is not allowed .

What is neutralisation?

What do you think would happen if we mixed an acid and a base?

Get learners to discuss this in class or in small groups. Allow them to speculate and guide them to recall their Grade 7 learning: An acid will lose its potency when it is mixed with a base and vice versa. So the acid will be weakened by the base and the base will be weakened by the acid. 'Weaken', however, is a term best avoided, because 'weak' and 'strong' have very specific meanings when speaking about acids and bases. In a sense their acid-base properties will be destroyed, because they will be converted to products that won't be acids or bases. (Often the salt that results from the reaction between an acid and a base will have acid-base properties of its own, but we will not be discussing that now.)

We are going to do an investigation to find out. We are going to mix vinegar with baking soda. But first, a little revision: is vinegar an acid or a base? If you are not sure, imagine putting a drop of vinegar on your tongue. What would it taste like?

It would taste sour, therefore it is an acid.

Is baking soda an acid or a base? If you are not sure, turn back to the previous chapter and look at the activity ' The pH scale'.

Baking soda is a base.

The reaction between vinegar and baking soda

Quantities for this investigation are as follows: Every 1 g of baking soda will require approximately 15 ml of vinegar for complete neutralisation. We recommend that you measure out 1 teaspoon of baking soda and approximately 50 ml vinegar for each group.

The purpose of this experiment is to investigate how the pH changes when vinegar is added to baking soda.

INVESTIGATIVE QUESTION(S):

A possible answer is: What will happen to the pH of the solution if we add vinegar to baking soda?

OVERVIEW OF THE INVESTIGATION:

In the range pH > 7

The pH will decrease.

HYPOTHESIS:

When we add vinegar to baking soda, the pH of the mixture will decrease.

• baking soda
• glass beaker or small yoghurt tub
• universal indicator paper (cut into 1 cm strips)
• sheet of white printer paper
• plastic teaspoon
• Prepare the universal indicator paper by neatly placing five 1-cm pieces underneath each other on the sheet of paper.
• Place one teaspoon of baking soda in the beaker or yoghurt tub.
• Add approximately 10 teaspoons of water to the baking soda.
• Use the teaspoon to stir the solution until all the baking soda has dissolved. We will be calling this the 'test solution' from now on.
• Transfer one drop of the test solution to the first piece of universal indicator paper using the teaspoon or a straw.
• Compare the colour of the paper with the colour guide given at the start of the chapter, to find the pH of the solution. Record this pH in your results table.
• Add 1 teaspoon of the vinegar to the test solution. Stir it gently and transfer another drop of the solution to a fresh strip of the universal indicator.
• Read the pH of the solution off the colour guide and record it in your results table.
• Repeat steps 6 and 7 until the pH of the test solution drops below 7. You may need more than 5 pieces of universal indicator paper.

Present your results in a neat table. Use appropriate headings for your table. 'Number of teaspoons of vinegar added' and 'Colour of the universal indicator paper' and 'pH of the test solution' are suggested headings for your columns.

Draw a line graph to illustrate your results. What will be on the x-axis and what will be on the y-axis? Give your graph a heading.

Learners must draw a graph with the 'number of teaspoons of vinegar added' on the x-axis (independent variable) and the pH of the solution on the y-axis (dependent variable).

CONCLUSIONS:

What conclusions can be made from the results of your investigation? Here you can rewrite your hypothesis, but change it to reflect your findings if they are different from what you predicted earlier.

Were you able to confirm or reject your hypothesis?

In this investigation, you probably noticed that the pH of the mixture dropped every time you added more vinegar to the baking soda! Why did this happen?

When an acid and a base are mixed (in the right amounts), they will neutralise each other. That means that, together, they will change into something that is neither an acid nor a base. So, the acid will lose its 'acidity' and the base will lose its 'basicity'.

What have we learnt so far? We have learnt that acids and bases neutralise each other:

If we add a base to an acid, the pH of the resulting solution will increase, because the acid will lose some of its potency .

• If we add an acid to a base, the opposite will happen. The pH will decrease, because the base will lose some of its potency.

What are the products of an acid-base reaction? Can we predict what they will be?

The products of acid-base reactions

In order to understand how an acid-base reaction works, we have to take a quick detour and say something about exchange reactions . Acid-base reactions are exchange reactions.

In the reaction below, two substances AB and CD are undergoing an exchange reaction:

AB + CD → AD + CB

Can you see that A and C have exchanged partners so that A is now combined with D, while C combined with B?

What does this have to do with acids and bases? Well, acids and bases undergo exchange reactions too. Here are some examples. See if you can figure out which parts have exchanged with which.

HCl + NaOH → NaCl + HOH

In the above equation HOH should actually be written: H 2 O. The reaction equation becomes:

HCl + NaOH → NaCl + H 2 O

or, in words:

hydrochloric acid + sodium hydroxide → sodium chloride + water

In this example, the following happened:

• the acid gave its H towards making a water molecule;
• the base gave OH towards making a water molecule; and
• the Na from the base and the Cl from the acid combined to form a salt.

2 HCl + MgO → MgCl 2 + HOH

2 HCl + MgO → MgCl 2 + H 2 O

hydrochloric acid + magnesium oxide → magnesium chloride + water

• the acid gave 2 H towards making a water molecule;
• the Mg from the base and the 2 Cl from the acid combined to form a salt.

Acid-base reactions always produce water and a salt. In both of the examples above the general equation was:

acid + base → salt + water

There is one class of acid-base reactions that produces an additional product, but we will learn more about that later.

In Grade 11 you will learn that the mechanisms of these reactions are actually slightly more complex than this, but for now, understanding it at this level is good enough.

Which laboratory acids should we know about?

When we investigated acids and bases in the previous chapter, we considered only household acids like lemon juice and vinegar. There are a few laboratory acids that you should know the names and formulae of and they have been listed in the following table:

These acids are very corrosive , even when they have been diluted with water and should always be handled with great care.

What happens if you put a burger in concentrated hydrochloric acid? (video)

In the next sections will discuss the classes of substances that are typically acids or bases. Two important things to remember are the following:

• Non-metal oxides form acidic solutions when they are dissolved in water.
• Metal oxides, metal hydroxides and metal carbonates all form basic solutions when they are dissolved in water.

First, we will look at the non-metal oxides.

Non-metal oxides form acidic solutions

Can you name a few non-metal oxides? Write down their formulae. If you are not sure you can take a peek at the Periodic Table and pick a few non-metals from the right-hand side of the table. Add oxygen and you have a non-metal oxide!

CO 2 and SO 2

How do we know that non-metal oxides form acidic solutions? Experiments have shown this.

You may not know this, but when CO 2 gas is bubbled through water some of it dissolves in the water to form carbonic acid. Here is the reaction equation:

CO 2 + H 2 O → H 2 CO 3

To see this happen, try the following quick activity.

CO 2 bubbled through water

A variation of this activity is if you have liquid universal indicator, you can add it to the tap water at the start to show the pH is 7 (it should be green). Then as you blow into the water, the universal indicator will change colour. Make sure you use a clear glass so learners can observe the colour change as it becomes more yellow. This links to the next activity on acid rain and how it forms.

• indicator paper

INSTRUCTIONS:

We could measure the pH of the water with universal indicator paper.

Now exhale into the water using a straw. Your breath contains CO 2 and some of this will dissolve in the water if you carry on doing this for a few minutes.

The pH will be below 7.

Carbonic acid is added to soft drinks to make it fizzy. The carbonic acid decomposes and forms carbon dioxide (CO 2 )

The pH of the solution is now below 7 because it contains carbonic acid (H 2 CO 3 ). Carbonic acid is not a very strong acid, but still acidic enough to have a pH lower than 7.

When sulfur dioxide (a gas) is bubbled through water it dissolves in the water to form an acid called sulfurous acid:

SO 2 + H 2 O → H 2 SO 3

These are two of the reactions that produce a phenomenon called acid rain . SO 2 and CO 2 are released as waste products from factories and power stations. For example, burning wood and fossil fuels releases carbon dioxide and sulfur dioxide into the atmosphere. These gases then dissolve in water droplets in the atmosphere to form acids, in a similar way that the CO 2 in your breath dissolved in the water in the last activity to produce an acidic solution. When it rains, these acids are present in the raindrops that fall back to earth. Sulfurous acid (H 2 SO 3 ) is strong enough to damage plant life and to acidify water sources.

For the next activity, you have to do some research on acid rain.

Volcanoes also release non-metal oxides into the air (mainly SO 2 ) that can contribute to acid rain.

What is acid rain?

• Study the diagram showing how acid rain forms.
• Do some extra reading and research about acid rain.
• Answer the questions about acid rain.

They are sulfur dioxide (SO 2 ), carbon dioxide (CO 2 ) and nitrogen dioxide (NO 2 ).

The main sources of these gases which contribute to acid rain are from human activity, such as electricity generation in fossil fuel power plants (especially coal), factories emitting smoke and the exhaust fumes from motor vehicles. Acid rain can also occur due to natural phenomena, such as volcanoes which emit sulfur dioxide into the atmosphere. Some processes in the ocean and in wetlands also produce the gases which form acids.

Sulphurous acid and carbonic acid.

The impacts include:

• damage of plant life, both wilderness areas and also crops, depending on where the rain falls
• the rain leaches into soil and makes it more acidic; this kills microorganisms living in the soil, damages plants further by contaminating soil water
• the rain can fall into various water sources and also run off into water sources such as rivers, lakes and dams; this causes the water to become more acidic; aquatic animals and plants can die; human water sources become too acidic as well

Acids are corrosive and so they can corrode surfaces over time.

Learners need to justify their answers. They may say that it helps the local environment as the gases are carried further away and therefore do not pollute the town or city that the factory is in or near. But this does not do anything to minimize the acid rain that could potentially form as the same amount of gases are still emitted; they are just carried further away. The acid rain therefore can still form and fall on the vegetation and areas outside of the towns and cities.

There are several solutions to minimizing the formation of acid rain. For example, coal-powered stations can use filters and other processes in their smoke towers to remove sulfur gases before the smoke is released into the atmosphere. Countries can take bigger steps by signing treaties to reduce their sulfur and other greenhouse gas emissions. The move towards using renewable energy sources will also help to reduce the reliance on coal and other fossil fuels, thereby reducing the emission of acid-producing gases into the atmosphere.

We have now learnt about non-metal oxides, but what about metal oxides? What kind of solutions do they form in water? We will find out more about them and other metal compounds in the next section.

Metal oxides, metal hydroxides and metal carbonates form basic solutions

Metal oxides.

Do you remember learning about some of the metal oxides in Chapter 3? We already learnt these rules to write the formulae of metal oxides.

Metal oxides from group 1 on the Periodic Table will have the formula M 2 O, where M represents any metal.

Can you write two examples? Look at the Periodic Table at the front of the book, pick any two metals from group 1 and write their formulae using this rule.

Any two of the following: Li 2 O, Na 2 O, K 2 O, Rb 2 O, Cs 2 O

Metal oxides from group 2 will have the formula MO.

Can you write 2 examples?

Any two of the following: BeO, MgO, CaO, SeO, BaO

What do you think the pH will be of a solution of a metal oxide in water?

The pH will be above 7.

The next class of compounds that form basic solutions in water are the metal hydroxides.

Metal hydroxides

A metal hydroxide forms when a metal reacts with water. A metal hydroxide has the general formula MOH or M(OH) 2 . In the formula, M represents a metal atom, O represents an oxygen atom and H represents a hydrogen atom.

To know whether the MOH or M(OH) 2 will be the correct formula, here are two simple rules for you to remember:

Metal hydroxides from group 1 on the Periodic Table will have the formula MOH.

Any two of the following: LiOH, NaOH, KOH, RbOH, CsOH.

Metal hydroxides from group 2 will have the formula M(OH) 2.

Can you write two examples?

Any two of the following: Be(OH)2, Mg(OH) 2 , Ca(OH) 2 , Sr(OH)2, Ba(OH)2.

What do you think the pH will be of a solution of a metal hydroxide in water?

The final class of compounds that forms basic solutions in water is the metal carbonates. Baking soda is a special kind of carbonate, called a bicarbonate (or hydrogen carbonate). You may remember that it was one of the bases we tested with universal indicator earlier.

Metal carbonates

A metal carbonate has the general formula MCO 3 or M 2 CO 3 . In the formula, M represents a metal atom, C represents a carbon atom and O represents an oxygen atom.

To know whether the MCO 3 or M 2 CO 3 will be the correct formula, there are two simple rules to remember:

Metal carbonates from group 1 on the Periodic Table will have the formula M 2 CO 3 .

Any two of the following: Li 2 CO 3 , Na2CO 3 , K 2 CO 3 ,Rb2CO 3 , Cs2CO 3 .

Metal hydroxides from group 2 will have the formula MCO 3 .

Any two of the following: BeCO 3 , MgCO 3 , CaCO 3 , SrCO 3 , BaCO 3 .

What do you think the pH will be of a solution of a metal carbonate in water?

The pH will be above 7

In the next sections we will be investigating real reactions!

The general reaction of an acid with a metal oxide

• metal oxide

In the previous section we learnt about two classes of oxides, namely metal oxides and non-metal oxides. Here is what we know about them so far:

• Metal oxides are formed from the reaction between a metal and oxygen. Metal oxides are basic. When we dissolve them in water, they form solutions with pH values above 7.
• Non-metal oxides are formed from the reaction between a non-metal and oxygen. Non-metal oxides are acidic. When they dissolve in water, they form solutions with pH values below 7.

Here is the same summary, in table form, with some examples added:

In this section, we are going to learn about the reactions between metal oxides and acids.

The reaction between magnesium oxide and hydrochloric acid

This investigation requires magnesium oxide from the reaction when magnesium ribbon burns in oxygen. If you have set some aside from the earlier activity ' The reaction of magnesium with oxygen' (Chapter 3), learners can use it for this investigation. If you did not, you can easily repeat that demonstration to produce more white magnesium oxide powder for this next investigation. This investigation is also suitable to scale up as a demonstration.

The purpose of this investigation is to:

• test whether a solution of magnesium oxide in water is acidic, basic or neutral; and

determine whether the reaction between an aqueous solution of magnesium oxide and hydrochloric acid is a neutralisation reaction .

What are the questions you hope to answer with this investigation? Write them in the space below. There are a few words to start you off.

When magnesium oxide is dissolved in water, will the resulting solution be acidic, basic or neutral?

When a solution of magnesium oxide is treated with hydrochloric acid, will the pH of the mixture increase, decrease, or stay the same?

The pH will decrease

What are your predictions? Your hypothesis should be a prediction of the finding(s) of the investigation. You should write it in the form of a possible answer to the investigative question(s). Here are a few words to start you off:

When magnesium oxide is dissolved in water, the resulting solution will be basic (have a pH > 7).

When a solution of magnesium oxide is treated with hydrochloric acid, the pH of the mixture will decrease.

• magnesium oxide powder
• white tile or sheet of white printer paper
• glass rod (or plastic straw)
• hydrochloric acid (HCl) solution (0.1 M)

Notes for the investigation:

Learners must not dilute hydrochloric acid themselves as it reacts strongly with water. Make sure to add the acid slowly to the water and NOT the other way around.

• To prepare 0.1 M HCl solution, carefully add approximately 10 ml concentrated hydrochloric acid (33% or 11 M) to 1 liter of tap water. It is recommended that you wear safety goggles and protective gloves during this step and that you rinse away any acid spills with cold tap water. Since this is just a qualitative experiment, it is not necessary to use distilled water for the solution. It is also not required that you measure the volumes with extreme accuracy.
• The following guide will help you to determine quantities: The magnesium oxide prepared from a 1 cm length of magnesium ribbon will require approximately 8 ml of 0.1 M hydrochloric acid for complete neutralisation. If the learners work in small groups and each group dissolves a small quantity of MgO (the size of a match head) in 2 ml of water, they will only need a few drops of HCl solution to neutralise all the MgO.

If you have universal indicator solution , this will work very nicely as you can observe the colour changes as you add the drops.

If you decide to give the learners droppers to measure out the HCl, you will have to enforce very strict rules for handling the droppers. Learners find the temptation to squirt water at each other very difficult to resist and they must be made aware of the hazards of accidentally squirting acid at another learner.

• Remind learners to use the colour guide for universal indicator provided at the start of the chapter. If you have the budget, a good idea would be to make a number of colour photocopies of the chart and to have them laminated so they will last longer.
• Remind learners to prepare a table for their results beforehand.
• Prepare the universal indicator paper by neatly placing five 1 cm pieces in a column on the white tile or sheet of printer paper.
• Place a small quantity (the size of a match head) of the magnesium oxide in a test tube.
• Add approximately 2 ml of tap water to dissolve most of the magnesium oxide.

Use the glass rod (or plastic straw) to stir the solution until most the magnesium oxide has dissolved. We will be calling this the test solution from now on.

• Transfer one drop of the test solution to the first piece of universal indicator paper.
• Compare the colour of the paper with the colour guide to find the pH of the solution.
• Record this pH in the table you prepared beforehand.
• Add 10 drops of the hydrochloric acid solution to the test solution. Stir it gently and transfer another drop of the solution to a fresh strip of the universal indicator.
• Read the pH of the solution off the colour guide and record it in your table.
• Repeat steps 3 and 4 until the pH of the test solution drops below 7. You may need more than 5 pieces of universal indicator paper.

Learner-dependent answer

Draw a graph of your results. Here are some hints to help you decide which variable to put on which axis:

Number of drops of HCl

The number of drops of HCl should be on the x-axis of the graph and pH should be on the y-axis. There should be a general trend downwards (since acid is added to a base, we can expect the pH to drop), but it should not be linear. This experiment is a very rudimentary 'titration' and an example of a titration curve from this experiment is given here:

It is therefore not expected that learners' curves will be linear, but rather that there will be a gradual decline in pH at first, followed by a rapid drop when all the base has been neutralised. After this the curve levels out again.

What conclusions can be made from the results of your investigation? Rewrite your hypothesis, but change it to reflect your findings if they are different from what you predicted earlier.

Were you able to confirm or reject your hypotheses?

Now that we have investigated a reaction between a metal oxide (MgO) and an acid (HCl), we can write an equation for the reaction. We will begin by writing a general equation and end with one that matches the reaction that we have just investigated.

A video showing the reaction of copper(II)oxide with hydrochloric acid

General equation for the reaction of an acid with a metal oxide

Can you remember learning that an acid-base reaction is an exchange reaction? We learnt that:

• The acid contributes H towards making a water molecule;
• The base contributes O or OH towards making a water molecule; and

Whatever is left of the the acid and the base after making a H 2 O molecule, combines to form a salt.

The general word equation for the reaction between an acid and a base is:

Since the base in our reaction is a metal oxide we can write:

acid + metal oxide → salt + water

This is the general word equation for the reaction between an acid and a metal oxide. The type of salt that forms will depend on the specific acid and metal oxide which were used in the reaction.

Equations for the reaction between magnesium oxide and hydrochloric acid

Now we are going to learn how to write equations for our actual reaction.

Writing the chemical equation

The following steps will guide you:

Magnesium oxide (MgO)

Now,let's try to predict the products of the reaction. We know that water will be one of the products.

Write what remains of the acid (HCl) after we have taken away the H (to make water). (Remember we need two H to make one H 2 O).

2 Cl (we used 2 HCl)

Now, let's put it all together, first the reactants, then the products:

Let's check quickly if the reaction is balanced.

2 H atoms left and 2 H atoms right. The H's are balanced.

2 Cl atoms left and 2 Cl atoms right. The Cl's are balanced.

1 O atoms left and 1 O atoms right. The O's are balanced.

Since the numbers of each type of atom is the same on either side of the equation, we can confirm that it is balanced.

Finally, let's use the chemical equation to write a word equation for the reaction:

In the next section we are going to look at the reactions between acids and metal hydroxides.

The general reaction of an acid with a metal hydroxide

• metal hydroxide

We will start this section with an investigation to illustrate the reaction between an acid and a metal hydroxide .

The reaction between sodium hydroxide and hydrochloric acid

• The same cautions regarding droppers and syringes apply to this activity. You will need to enforce very strict rules for handling these items or learners may find the temptation to squirt water at each other very difficult to resist.
• Remember to provide learners with a colour guide for universal indicator, if you have this.
• Remind learners to draw a results table before they start the experiment.
• test whether sodium hydroxide is acidic or basic; and
• determine whether the reaction between sodium hydroxide and hydrochloric acid is a neutralisation reaction.

Here are some ideas:

• When sodium hydroxide is dissolved in water, will the resulting solution be acidic, basic or neutral?
• When a solution of sodium hydroxide is treated with hydrochloric acid, will the pH of the mixture increase, decrease or stay the same?
• Will it be possible to neutralise all the sodium hydroxide by adding hydrochloric acid?

OVERVIEW OF THE INVESTIGATION :

Some ideas:

• Sodium hydroxide solution will have a pH greater than 7.
• When a solution of sodium hydroxide is treated with hydrochloric acid, the pH of the mixture will decrease.
• By adding hydrochloric acid to the sodium hydroxide solution, it should be able to decrease the pH to 7 and even below 7.
• sodium hydroxide solution (0.1 M)

Prepare 0.1 M NaOH solution by dissolving approximately 4 g of NaOH pellets in 1 liter of cold tap water. Wear safety goggles and gloves since there is a chance the sodium hydroxide solution could splash up.

• glass rod or plastic straw
• test tube or small glass beaker
• plastic syringe (2.5 ml capacity)

Instructions for preparation are given with the previous investigation: The reaction between magnesium oxide and hydrochloric acid

Use the syringe to transfer 2 ml of the sodium hydroxide solution into the test tube or small glass beaker. We will be calling this the test solution from now on.

• Rinse the syringe very thoroughly with water and dry it out with a clean tissue. Now fill it with hydrochloric acid solution and set it aside.
• Transfer one drop of the sodium hydroxide (test solution) to the first piece of universal indicator paper.
• Compare the colour of the paper with the colour guide to find the pH of the sodium hydroxide solution. Record this pH in your results table.
• Add 0.5 ml of the hydrochloric acid solution from the syringe to the test solution. Stir it gently with the glass rod or straw and transfer another drop of the test solution to a fresh strip of the universal indicator paper.

Learner dependent answer. Should be around 2 ml.

• If you are quite sure that all the base has been neutralised by the acid (the pH should be 7 and the universal indicator paper should have turned green), pour the test solution into a small glass beaker and leave it in the window sill for a few days. Remember to come back to it later to see what has happened to it.

NaCl forms in this reaction and the idea is for learners to let it dry out in the window sill and examine it later. It is probably not a good idea to let them taste it, as there is a possibility that not all of the acid or base has been neutralised.

learner-dependent answer

Draw a graph of your results.

Volume of HCl

The volume of HCl added should be on the x-axis of the graph and pH should be on the y-axis. There should be a general trend downwards (since acid is added to a base, we can expect the pH to drop), but it should not be linear. See comments and graph provided with the previous investigation.

Now that we have investigated a reaction between a metal hydroxide (NaOH) and an acid (HCl), we can write an equation for the reaction. We will begin by writing a general equation and end with one that matches the reaction that we have just investigated.

A video illustrating the reaction of magnesium hydroxide with hydrochloric acid in the presence of universal indicator

General equation for the reaction of an acid with a metal hydroxide

You learnt that an acid-base reaction can be represented by the following general word equation:

The base in our reaction was a metal hydroxide, so the general equation becomes:

acid + metal hydroxide → salt + water

This is the general equation for the reaction between an acid and a metal hydroxide. The type of salt that forms will depend on the specific acid and metal hydroxide which were used in the reaction.

Equations for the reaction between sodium hydroxide and hydrochloric acid

Sodium hydroxide (NaOH)

Write what remains of the acid after we have taken away the H to make water. Remember we need two H to make one H 2 O, but NaOH has already contributed one O and one H. Now put the two fragments together. Place the metal from the base first and the non-metal from the acid. One Na and one Cl makes...

Now, let's put it all together, in the following order: Acid + metal hydroxide → salt + water

2 H atoms on the left and 2 H atoms on the right. The H's are balanced.

1 Cl atoms on the left and 1 Cl atoms on the right. The Cl's are balanced.

1 O atoms on the left and 1 O atoms on the right. The O's are balanced.

Once you have performed this reaction and you are left with a neutral solution , you decide you want to recover the sodium chloride (table salt). How will you do this?

You need to evaporate the water so that the salt crystallizes, either by leaving it in a sunny spot or boiling the solution.

In the next section we are going to look at the reactions between acids and metal carbonates.

The general reaction of an acid with a metal carbonate

• metal carbonate

In this section we will investigate the reaction between an acid and a metal carbonate .

The reaction between calcium carbonate (chalk) and hydrochloric acid

Grind up a few pieces of white chalk for this experiment. The calcium carbonate will not actually dissolve well in water, but it should be possible to determine that the solution is basic, from the tiny amount of calcium carbonate that will dissolve when the chalk dust is suspended in water.

Learners will need their colour charts and results tables before they start.

• test whether calcium carbonate is acidic or basic;
• determine whether the reaction between calcium carbonate and hydrochloric acid is a neutralisation reaction; and
• determine the products of the reaction between calcium carbonate and hydrochloric acid.

INVESTIGATIVE QUESTIONS:

• When calcium carbonate is dissolved in water, will the resulting solution be acidic, basic or neutral?
• When a solution of calcium carbonate is treated with hydrochloric acid, will the pH of the mixture increase, decrease or stay the same?
• Will it be possible to neutralise all the calcium carbonate by adding hydrochloric acid? (Be careful not to introduce misconceptions here)
• What other products will form when calcium carbonate reacts with hydrochloric acid?

We will measure the pH of a suspension of calcium carbonate (CaCO 3 ) with universal indicator paper to confirm whether it is acidic or basic. Within what range do you expect the pH of the calcium carbonate to fall?

• Calcium carbonate solution will have a pH greater than 7
• When calcium carbonate is treated with hydrochloric acid, the pH of the mixture will decrease
• By adding hydrochloric acid to the calcium carbonate, it should be able to decrease the pH to 7 and even below 7
• chalk dust (calcium carbonate) suspended in a small quantity of water.
• plastic syringe (2.5 cm capacity) or dropper

Place approximately 2 ml of the calcium carbonate suspension into the test tube or small glass beaker. We will be calling this the test solution from now on.

• Transfer one drop of the calcium carbonate (test solution) to the first piece of universal indicator paper.
• Compare the colour of the paper with the colour guide below, to find the pH of the calcium carbonate solution. Record this pH in your results table.
• Add 0.5 cm of the hydrochloric acid solution from the syringe to the test solution. Watch very carefully what happens. Do you see anything interesting? (Hint: Look for bubbles!) Stir the test solution gently with the glass rod and transfer another drop of it to a fresh strip of the universal indicator.
• Repeat steps 6 and 7 until the pH of the test solution reaches approximately 7. How much of the hydrochloric acid solution have you used? Write the volume in the space below.
• Your teacher will repeat the experiment as a demonstration and will collect the gas that formed during the reaction, for testing with clear limewater.

Perform the experiment in a conical flask, as follows:

2 into another conical flask containing lime water (see diagram below). The CO 2 gas should be poured out. Shake the conical flask containing the lime water and CO 2 to facilitate mixing. Allow the learners to make their observations.2 is denser than air and will remain in the conical flask for a few minutes before diffusing into the air. It is during this time that you should pour it over into the lime water. Be careful not to let any of the test solution flow over into the clear lime water. Only the CO

Alternatively, you could use a setup like the one shown in the diagram below:

These experiments can also be done using combo plates.

Carbon dioxide, CO 2

Now that we have investigated a reaction between a metal carbonate (CaCO 3 ) and an acid (HCl), we can write an equation for the reaction. We will begin by writing a general equation and end with one that matches the reaction that we have just investigated.

A video showing the reaction between metal carbonates and acids

General equation for the reaction of an acid with a metal carbonate

The general equation for the reaction between an acid and a base is as follows:

If we replace 'base' with 'metal carbonate', we get:

acid + metal carbonate → salt + water

But wait, there was a third product in our reaction! Can you remember what it was? (Hint: Bubbles formed, so it was a gas.)

We need to make it clear that CO 2 was a product of the reaction, so the correct general word equation would be:

acid + metal carbonate → salt + water + carbon dioxide

The type of salt that forms will depend on the specific acid and metal carbonate which were used in the reaction.

Equations for the reaction between calcium carbonate and hydrochloric acid

Calcium carbonate (CaCO 3 )

Write what remains of the base after we have taken away the CO 3 to make CO 2 and leave one O to make water.

Write what remains of the acid after we have taken away the H to make water. Remember we need two H to make one H 2 O and CaCO 3 has only contributed one O.

2 HCl are needed, so 2 Cl will remain.

2 HCl + CaCO 3 → CaCl 2 + H 2 O + CO 2

2 H left and 2 H right. The H's are balanced.

2 Cl left and 2 Cl right. The Cl's are balanced.

3 O left and 3 O right. The O's are balanced.

1 C left and 1 C right. The C's are balanced.

hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide

Applications for calcium carbonate

Calcium carbonate is found in many places outside of the laboratory. It is found in different types of rocks around the world, for example limestone, chalk and marble.

Calcium carbonate is also the main part of shells of various marine organisms, snails, pearls, oysters and bird eggshells. It is also found in the exoskeletons of crustaceans (such as crabs, prawns and lobsters).

Dark green leafy vegetables such as broccoli, kale and cabbage are a dietary source of calcium carbonate, providing the body with calcium. You can also take calcium carbonate in the form of tablet supplements.

Calcium carbonate also has many applications. In industry, the main application is in construction as it is used in various building materials and in cement. Calcium carbonate is used in many adhesives, paints and in ceramics. It is also used in swimming pools to adjust the pH. When do you think it would be added? If the pool was too acidic and you wanted to make it more basic, or if the pool was too basic and you wanted to make it more acidic?

CaCO 3 forms a basic solution in water so it is used if the pH is too low (too acidic) and you want to make the pool water more basic.

Calcium carbonate is also used in agriculture in the form of lime powder. Agricultural lime is made by grinding up limestone or chalk. It is added to the soil if the soil is too acidic to increase the pH. It also provides plants with a source of calcium.

In this chapter we have investigated a number of reactions of acids with bases. We have learnt to write word equations for these reactions and practised converting between word and balanced chemical equations.

• The reaction of an acid with a base is called a neutralisation reaction.
• When an acid (pH 7), the pH of the resulting mixture will lie somewhere between that of the acid and the base. Even though the acid and base will be neutralised, the resulting solution will not necessarily be neutral.

Some common laboratory acides are sulfuric acid (H 2 SO 4 ), nitric acid (HNO 3 ) and hydrochloric acid (HCl).

• Non-metal oxides tend to form acidic solutions when they dissolve in water. These solutions will have pH values below 7.
• Metal oxides, metal hydroxides and metal carbonates form basic solutions in water; these will have pH values above 7.
• When a metal oxide, or a metal hydroxide reacts with an acid, a salt and water form as products.
• When a metal carbonate reacts with an acid, a salt, water and carbon dioxide form as products.

Concept map

Complete the concept map by filling in the blank spaces..

This is the completed concept map.

Revision Questions

Fill in the missing words in these sentences. Write the word on the line below. [10 marks]

To know if something is an acid or a base, we measure its _______.

The name of the laboratory acid with the formula H 2 SO 4 , is _______.

sulfuric acid

The formula of the laboratory acid named hydrochloric acid, is _______.

When a metal oxide reacts with an ______, a salt and water will be formed.

When a metal hydroxide reacts with an acid, a salt and _____ will be formed.

When a metal carbonate reacts with an acid, a salt, water and _______ will be formed.

carbon dioxide

Metal oxides, metal hydroxides and metal carbonates all dissolve in water, forming ______ solutions. This means the solutions will have pH values ____ than 7.

basic; greater

The reaction of an acid with a base is called a _______ reaction.

neutralisation

Non-metal oxides tend to form ______ solutions when they dissolve in water.

Write a short paragraph (3 or more sentences) to explain what you understand each of the following terms to mean, in your own words. [2 x 3 = 6 marks]

Learner's paragraph should contain at least the following ideas:

• When an acid and a base are mixed, the acid will lose some of its 'acidity' and the base will lose some of its 'basicity'.

If they are mixed in the right amounts, they will neutralise each other.

• The products of the reaction will be a salt and water.
• Certain industries (and even some natural phenomena like volcanic eruptions) produce non-metal oxides as waste products.
• Non-metal oxides form acidic solutions when they dissolve in atmospheric water droplets.
• These acidic solutions rain down onto the Earth's surface and can cause damage to buildings, plant life and acidify water sources.

For each of the following reactions, complete the tables by providing the missing equations.

The reaction between hydrochloric acid and magnesium oxide [4 marks]

The reaction between hydrochloric acid and sodium hydroxide [6 marks]

The reaction between hydrochloric acid and calcium carbonate [4 marks]

The reaction between hydrochloric acid and magnesium hydroxide [4 marks]

The reaction between hydrochloric acid and calcium oxide [4 marks]

The reaction between hydrochloric acid and potassium hydroxide [6 marks]

Total [48 marks]

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Copper and hydrochloric acid

When hydrochloric acid is added to a sample of copper why does it look pink at first then form a pale blue solution ?

• experimental-chemistry

• 1 $\begingroup$ Well copper is pink... $\endgroup$ –  Mithoron Commented Sep 8, 2015 at 22:55
• 1 $\begingroup$ The copper was not pure but contaminated with another substance. Likely your sweaty oily hands. $\endgroup$ –  user5434678 Commented Oct 11, 2015 at 16:53

If pure hydrochloric acid is added to a sample of pure copper, there is essentially no reaction; the solution should not turn blue.

At a guess, you're seeing the copper turn pink because it has a significant oxide layer that is reacting with the acid, forming copper (II) chloride and leaving bare the pinkish metal. The solution could be pale blue or pale green, and the time delay of the color would be a result of the amount of time required for the reaction to occur. Note that if this is the case, you would not expect to see any bubbling on the copper, as copper (II) oxide will react with hydrochloric acid to form copper (II) chloride and water.

As Ivan Neretin pointed out in comments, large pieces of copper will react with hydrochloric acid over a period of several hours to days in the presence of air. Permeakra added that a fine copper powder will react in minutes with a solution of hydrochloric acid in the presence of air. In both cases it is the addition of oxygen that causes the copper to react, and in both cases the solution will be a greenish yellow color due to the excess of hydrochloric acid required.

As an example, a student of mine recently asked about this with regard to a penny. We cut open a modern US penny (95% zinc, 5% copper by mass) to expose the zinc core, and placed it in 200mL of 6M HCl. The oxidation layer on the penny was gone almost immediatley after it entered the HCl solution. The zinc reacted quickly and by the next day the penny had been converted into a hollow copper shell. The shell sat in the remaining HCl solution (still 5.7M) over the weekend with no visible reaction but the solution slowly turned green and the copper dissolved completely by the following Friday.

If your solution is turning pink, then I'm not entirely sure what's going on.

• $\begingroup$ In fact, copper would slowly dissolve in acid in the presence of air , even if initially there were no oxide layer at all. $\endgroup$ –  Ivan Neretin Commented Sep 9, 2015 at 6:11
• $\begingroup$ @IvanNeretin Indeed, but this takes a loooong time before the reaction is noticeable, at least with the acid concentrations I've used for this (several days.) Though kay_student doesn't specify a timeline, I assumed it was faster than that. I thought about mentioning the oxidizing acid aspect of HCl and air, but decided against it for simplicity's sake. $\endgroup$ –  Jason Patterson Commented Sep 10, 2015 at 0:50
• $\begingroup$ @JasonPatterson This depends on the copper sample. Fine powder in excess of acid dissolves within minutes. $\endgroup$ –  permeakra Commented Oct 11, 2015 at 17:11
• $\begingroup$ @permeakra Good point. Edited to include additional information from previous comments. $\endgroup$ –  Jason Patterson Commented Oct 11, 2015 at 17:13

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Generating, collecting and testing gases

In association with Nuffield Foundation

Gases give rise to particular hazards so great care must be taken when preparing, collecting or testing

Gases give rise to particular hazards so great care must be taken when preparing, collecting or testing them.

How the gas is to be used will differ from experiment to experiment – it is essential to read carefully the specific instructions given or referred to in the experimental procedure and any accompanying technical notes. This is especially important if the gas needs to be dried.

Gases can be collected by upward or downward delivery or over water. Refer to specific information on each gas below.

Health, safety and technical notes 

• Read our standard health and safety guidance.  
• Wear eye and skin protect if required.
• Hydrochloric acid is highly corrosive, refer to CLEAPSS Hazcards  HC047a , plus CLEAPSS Recipe Book RB021.  
• Zinc is of low hazard, refer to CLEAPSS Hazcard  HC108b
• Copper is of low hazard, refer to CLEAPSS Hazcard  HC026 .
• Hydrogen gas is EXTREMELY FLAMMABLE – ensure there are no naked flames. Refer to CLEAPSS Hazcard  HC048
• Hydrogen peroxide (IRRITANT) to manganese(IV) oxide powder (HARMFUL). Collect the gas over water. Refer to CLEAPSS Hazcard  HC050
• Oxygen is an OXIDISING AGENT, refer to CLEAPSS Hazcard  HC069
• Potassium manganate(VII) (OXIDISING, HARMFUL and DANGEROUS FOR THE ENVIRONMENT). Refer to CLEAPSS Hazcard  HC081

Double-check that the acid is hydrochloric and NOT sulfuric.

• Sodium chlorate(I) can cause headache, fatigue, dizziness, and methemoglobinemia. Refer to CLEAPSS Hazcards  HC089 .
• Chlorine is classified as toxic, irritant and dangerous for the environment. Refer to CLEAPSS Hazcard  HC022a

General gas preparation

The diagram below shows a typical set of apparatus which can be used to prepare a range of gases.

Source: Royal Society of Chemistry

Typical apparatus used for preparing a range of gases

Gas collection methods

The diagrams below show three different methods for collecting gas.

Three methods of gas collection

Preparing specific gases

Wear appropriate eye protection. The amounts given below are sufficient to generate 1 litre (1 dm 3 ) of each of the named gases.

Carbon dioxide, CO 2

Slowly add 42 cm 3  of 2 M hydrochloric acid (IRRITANT) to an excess of marble chips. Collect the gas by downward delivery or over water (slightly soluble).

Hydrogen, H 2

Slowly add 28 cm 3  of 3 M hydrochloric acid (CORROSIVE) to excess zinc granules and 1 g of hydrated copper sulfate (HARMFUL). Collect the gas by upward delivery or over water.

Hydrogen gas is EXTREMELY FLAMMABLE – ensure there are no naked flames.

Oxygen, O 2

Slowly add 50 cm 3  of 20 vol hydrogen peroxide (IRRITANT) to manganese(IV) oxide powder (HARMFUL). Collect the gas over water.

Oxygen is an OXIDISING AGENT.

Chlorine, Cl 2

Work in a fume cupboard. Method 2 is safer and recommended but slower.

Add 14 cm 3  of concentrated hydrochloric acid (CORROSIVE) to at least 3 g of potassium manganate(VII) (OXIDISING, HARMFUL and DANGEROUS FOR THE ENVIRONMENT).

Add 5 M hydrochloric acid (IRRITANT) to 30 cm 3  of recently purchased (10–14% available chlorine) sodium chlorate(I) solution (CORROSIVE) with plenty of stirring. Note that sodium chlorate(I) is only available as a solution often called ‘sodium hypochlorite’; it must not be confused with sodium chlorate(V) (sometimes just called ‘sodium chlorate’), which is a white, crystalline solid. School samples often react too slowly because old sodium chlorate(I) is used. This will have less than the required 10% available chlorine (as it applies to both methods). 

The equipment required for preparing chlorine gas

Collect the gas by downward delivery. Chlorine is classified as TOXIC, IRRITANT and DANGEROUS FOR THE ENVIRONMENT.

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

The resource is also part of the Royal Society of Chemistry’s Continuing Professional Development course:  Chemistry for non-specialists .

© Nuffield Foundation and the Royal Society of Chemistry

• Teacher notes
• Technician notes
• Practical skills and safety

Specification

• 2.9.5 describe the laboratory preparation and collection of hydrogen using zinc (or other suitable metal) and hydrochloric acid, and recall the physical properties of hydrogen and its uses, including weather balloons and hardening oils, and its potential…
• 2.9.6 describe the laboratory preparation and collection of oxygen by the catalytic decomposition of hydrogen peroxide, and recall the physical properties of oxygen and its uses in medicine and welding;
• 2.9.8 describe the laboratory preparation and collection of carbon dioxide gas using calcium carbonate and hydrochloric acid, and recall the uses of carbon dioxide in fizzy drinks and fire extinguishers.
• 7. Investigate the effect of a number of variables on the rate of chemical reactions including the production of common gases and biochemical reactions.
• Mandatory experiment 6.1 - Monitoring the rate of production of oxygen from hydrogen peroxide, using manganese dioxide as a catalyst.
• Demonstration of the effects on reaction rate of (i) particle size
• 4.8.2.1 Test for hydrogen
• 4.8.2.2 Test for oxygen
• 4.8.2.3 Test for carbon dioxide
• 4.8.2.4 Test for chlorine
• Describe tests to identify selected gases including hydrogen and carbon dioxide.
• Describe tests to identify selected gases including oxygen, hydrogen and chlorine.
• 5.8.2.1 Test for hydrogen
• 5.8.2.2 Test for oxygen
• 5.8.2.3 Test for carbon dioxide
• 5.8.2.4 Test for chlorine
• 3.12 Describe the chemical test for: hydrogen, carbon dioxide (using limewater)
• C1.1.12 describe tests to identify oxygen, hydrogen and carbon dioxide
• C1.4.2 describe a test to identify chlorine (using blue litmus paper)
• C4.2a describe tests to identify selected gases
• C1.3g describe tests to identify selected gases
• Simple tests can be used to identify oxygen, hydrogen and carbon dioxide gases.
• methods for the collection of gases including:
• (o) the reactions of the alkali metals with air/oxygen, the halogens and water
• (p) the test used to identify hydrogen gas
• (j) the tests used to identify oxygen gas and carbon dioxide gas
• (n) the reactions of the alkali metals with air/oxygen, the halogens and water
• (o) the test used to identify hydrogen gas
• (m) reactions of the halogens with metals

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By Dorothy Warren and Sandrine Bouchelkia

Practical experiment where learners produce ‘gold’ coins by electroplating a copper coin with zinc, includes follow-up worksheet

Practical potions microscale | 11–14 years

By Kirsty Patterson

Observe chemical changes in this microscale experiment with a spooky twist.

Antibacterial properties of the halogens | 14–18 years

By Kristy Turner

Use this practical to investigate how solutions of the halogens inhibit the growth of bacteria and which is most effective

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