CBSE Class 12 Chemistry Chemical Demonstration all chapters

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CHEMICAL DEMONSTRATIONS

TOPIC: PHYSICAL AND CHEMICAL CHANGES

Demonstration: Distinguishing between chemical and physical change

Concepts

• A physical change does not involve any bond breaking or making and no new chemical is formed; only the state of the chemical is changed, e.g. from solid to liquid.
• A chemical change involves the breaking and making of bonds and the formation of a new substance, identified by, e.g., a change in colour.

Safety

• Safety glasses must be worn; handle lead nitrate and potassium iodide with care.
• Minimise contact with skin by wearing gloves and using a spatula if possible when handling these chemicals.

Requirements

• Physical change: ice.
• Chemical change: 2 large test tubes, a rubber stopper to fit the test tube, 10 g lead nitrate, 10 g potassium iodide, a black background.

Preparation

Make sure that both lead nitrate and potassium iodide are in fine powder form; if not, grind each solid separately using a mortar and pestle prior to the demonstration.

Demonstration instructions

• Physical change: Using tongs, transfer some of the ice into a conical flask. Shake the flask to increase the rate of change from solid to liquid.
• Chemical change: (1) Place about 15 g of lead nitrate into one test tube and about 15 g of potassium iodide into another test tube, then transfer one solid into the other. (2) Stopper the test tube and shake vigorously until there is a colour change.

Observations

• Physical change: Ice changes into liquid, i.e. water.
• Chemical change: Both lead nitrate and potassium iodide are white solids. Upon mixing, a yellow solid (lead iodide) is produced. A change in colour indicates a chemical change rather than a physical change. Mixing two solids to produce a chemical reaction is unusual, as normally this reaction is done by mixing lead nitrate solution and potassium iodide solution; however, using solids is more effective as a demonstration.
• Students will observe that in the first instance there was only a physical change, since solid H₂O turned into liquid, whereas in the second example a new substance with a different colour was formed with no white powder remaining, which means that both potassium iodide and lead nitrate have converted to a new substance.

Learning outcomes

• Students will be able to distinguish between physical and chemical changes.
• Students will be able to write balanced equations from observations and include physical states for all reactants and products.

Chemical disposal

Solid waste should be disposed of into a special toxic waste bin.

Reaction

Pb(NO₃)₂ + 2KI → PbI₂ + K₂(NO₃)₂

TOPIC: CHEMICAL REACTIONS

Demonstration: Burning magnesium

Concepts

• Characteristics of a chemical change
• Identification of products of a reaction
• Balancing chemical equations

Safety

• Safety glasses MUST be worn at all times.
• DO NOT look directly at the light produced by the magnesium.
• Work on a heat-proof mat when using open flames.
• Turn off the gas/spirit lamp when not in use.

Requirements

Safety glasses, heat-proof mat, metal tongs, gas burner/spirit lamp, matches, magnesium ribbon strips, phenolphthalein.

Preparation

Take small strips of magnesium and clean off (use sand paper) any oxide layer formed on the magnesium prior to burning.

Demonstration instructions

1. Put the spirit lamp on top of the heat-proof mat. Turn on the spirit lamp.
2. Using metal tongs, pick up a piece of magnesium ribbon and hold it inside the hottest part of the flame. Wait until the magnesium starts to burn, then take it out of the flame. Please make sure you tell students NOT to look directly at the light produced by the magnesium.
3. Drop the burning magnesium ribbon onto the heat-proof mat.
4. Transfer the burnt magnesium into water, shake the contents, and heat the contents.
5. Add a few drops of phenolphthalein to the contents.

Observations

1. The heated magnesium strip produces a bright white light while burning.
2. The magnesium changes from a silver metallic strip to a white powder.
3. The burnt magnesium is magnesium oxide.
4. Magnesium oxide is sparingly insoluble in cold water but fairly soluble on hot water.
5. The aqueous solution of magnesium oxide is basic; that is why it changes the colour of phenolphthalein from colourless to pink.

Chemical equation

2Mg(s) + O₂(g) → 2MgO(s)

MgO(s) + H₂O(l) → Mg(OH)₂(aq)

Learning outcomes

• Students should be able to convert observations into a chemical equation.
• Students should be able to balance chemical equations.
• Students should identify states of matter and include those when writing equations.

Disposal of chemicals

All solids may be placed into a laboratory bin.

Demonstration: Reaction between sodium and water

Concepts

• Characteristics of a chemical change
• Converting observations into a chemical equation
• Identifying products of a reaction
• Balancing chemical equations

Safety

• Safety glasses MUST be worn at all times.
• MUST use tweezers when handling sodium.
• Take only one piece of sodium at a time out of the paraffin oil.
• Observe from a distance. Do not stand too close while the reaction is taking place.
• Only use small pieces of sodium (no bigger than pea size).

Requirements

Safety glasses, overhead projector, large glass bowl, water, tweezers, scalpel, ceramic plate, sodium metal in a container with paraffin oil, phenolphthalein indicator.

Preparation

Cut three pea-size pieces of sodium metal and place them in paraffin oil in an enclosed container. Half-fill the large glass bowl with water.

Demonstration instructions

1. The bowl of water should be placed on top of the overhead projector.
2. Place about five drops of phenolphthalein indicator into the water.
3. Using tweezers, drop a small piece of sodium metal into the water.

Observations

• A vigorous reaction takes place upon dropping the sodium metal into the water.
• A pink colour is formed as the metal reacts with the water.
• A colourless gas is evolved which ignites and a small flame is produced (depends on the size of the sodium).
• As soon as the sodium is dropped into water, it starts to turn pink. The colour is due to the phenolphthalein indicator, which turns pink in the presence of hydroxide ions; as the reaction proceeds more hydroxide ions are formed and a pinker colour appears.

Chemical equation

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Learning outcomes

• Students should gain practice in converting observations from a reaction into a chemical equation.
• Students should be able to balance chemical equations and identify states of matter.
• Students should understand the purpose of an indicator in this reaction and hence why the water turns pink once the sodium metal is dropped in.

Disposal of chemicals

Leftover solution should be washed down a laboratory sink with plenty of water. Make sure any leftover unused sodium pieces are returned to the paraffin oil container. DO NOT place ANY sodium pieces (no matter how small) down the sink.

Demonstration: Reacting hydrogen gas with oxygen gas to produce water

Concepts

• Characteristics of a chemical change
• Converting observations into a chemical equation
• Identifying products of a reaction
• Balancing chemical equations

Safety

• Safety glasses must be worn at all times.
• Tie long hair back.
• Note that this reaction produces a large flame and some smoke. DO NOT do this reaction in a small classroom. The reaction may be done in a lecture theatre with plenty of overhead space or, alternatively, outdoors away from flammables.
• Turn off smoke sensors to avoid an unnecessary evacuation, but note the location of the manual fire alarm and have a fire extinguisher close by in case of a fire.
• Do not fill the balloon up too much with hydrogen, as this will lead to a larger explosion. A small balloon works just as well and is safer and not as loud.

Equipment

Safety glasses, a few balloons, hydrogen gas, several pieces of string (about 2 m each), a long stick with a wax taper secured to the end, matches.

Preparation

Blow up the balloons with hydrogen gas. If you are using the reaction between zinc and hydrochloric acid as your source of hydrogen gas, try not to let any air into the balloon, as this may cause the balloon not to explode but to melt and burn. Tie a long piece of string to the end of each balloon.

Demonstration instructions

1. Tie the end of the string around the leg of a chair or table. Make the string as long as possible, without allowing the balloon to touch the ceiling and so that you are able to reach it with the wax taper stick.
2. Light the wax taper.
3. Stand as far away as possible from the balloon and hold out the wax taper until it touches the balloon. WARN students of the bang.

Observations

• A very loud bang is heard.
• A large ball of flame is produced.
• Inspection of the remains of the balloon shows small droplets of water formed inside the balloon.

Chemical equation

2H₂(g) + O₂(g) → 2H₂O(l)

Learning outcomes

• Students should be able to write a balanced equation from observation and identify states of matter.
• Students should gain practice in converting observations from a reaction into a chemical equation.

Disposal of chemicals

Unexploded hydrogen balloons should be leaked out away from sparks.

TOPIC: PROPERTIES OF GASES

Demonstration: Preparation and properties of oxygen gas

Concepts

• Properties of oxygen gas
• Method used in identification of oxygen gas
• The role of a catalyst in a reaction
• Balancing chemical equations

Safety

• Safety glasses MUST be worn at all times.
• Take care when handling glass tubing, particularly when inserting it into stoppers or plastic tubing.
• Do not place a lighted wooden splint inside the O₂-containing test tube.

Requirements

• Safety glasses
• A large plastic container
• 1 L beaker
• 3 test tubes
• A delivery tube with a stopper on one end
• A clamp and stand
• Spatula
• 10 mL measuring cylinder
• Wooden splint and matches
• Two rubber stoppers
• Manganese(IV) dioxide
• 20 mL of 20% hydrogen peroxide

Preparation

Three-quarters fill a 1 L beaker with tap water. Fill two test tubes with water. Place a finger over the tube mouth, invert, and place the unsealed end under the surface of the water before removing your finger. The test tube should not contain any air. Place the end of the delivery tube into the beaker. Clamp the reaction tube.

Demonstration instructions

1. Do this reaction within the plastic container to contain any spills.
2. Place a small quantity (about a spatula full) of MnO₂ into a reaction tube and add 5 mL of 20% H₂O₂. Quickly replace the stopper assembly.
3. Allow a little gas to pass through the delivery tube (to displace the air) and then collect a test tube of oxygen by directing the tubing underneath one of the water-filled test tubes. If the reaction stops, add a further 5 mL of H₂O₂.
4. When the test tube is filled with oxygen, stopper it and remove it from the beaker of water.
5. Place a glowing, but not alight, wooden splint into the test tube.

Observations

• The reaction is quite vigorous, with the black MnO₂ foaming quickly.
• A colourless, odourless gas (O₂) is seen filling up the water-filled test tubes.
• Oxygen is collected by downward displacement of water.
• A glowing splint re-lights when placed inside the test tube.

Chemical equation

MnO₂ + 2H₂O₂(l) → O₂(g) + 2H₂O(l)

Learning outcomes

• Students should be able to explain properties of oxygen gas.
• Students should be able to explain the role of a catalyst.
• Students should be able to write a balanced chemical equation for the preparation of oxygen gas.

Disposal of chemicals

Liquid waste should be washed down the sink with plenty of water.

TOPIC: OXIDISING AGENTS AND REACTION RATES

Demonstration: Decomposition of hydrogen peroxide

Concepts

• Oxidising agents
• Decomposition reactions
• Reaction rates and catalysts

Safety

• Safety glasses MUST be worn at all times.
• Wear safety gloves when handling 30% hydrogen peroxide.

Requirements

100 mL of 30% hydrogen peroxide, a pea-size amount of potassium permanganate, 5 L conical flask, spatula, overhead projector.

Preparation

No prior preparation required.

Demonstration instructions

Add enough hydrogen peroxide to cover the bottom of the conical flask (about 30 mL into a 5 L flask). Place the flask on top of a projector and turn on the light. Then quickly add a pea-sized amount of KMnO₄ to the flask. Make sure you are wearing safety glasses.

Observations

• When the peroxide is poured into the flask, it looks like it is “boiling” as it slowly starts to decompose to a colourless gas.
• When potassium permanganate is added, suddenly all of the hydrogen peroxide is decomposed and, due to lots of heat being given off, steam is formed.
• At the end of the decomposition process, potassium permanganate is seen at the bottom of the flask.

Chemical equation

2H₂O₂(l) → 2H₂O(l) + O₂(g) (KMnO₄ as catalyst)

Learning outcomes

• Students should appreciate the role of a catalyst in increasing the rate of reaction and not taking part in the reaction.
• Students should be able to understand decomposition reactions.
• Students should be able to write a balanced chemical equation for the decomposition of hydrogen peroxide.

Disposal of chemicals

Rinse out the flask with water and wash it down the laboratory sink with plenty of water.

TOPIC: CHANGE OF STATE

Demonstration: Sublimation of iodine

Concepts

• Change of state from solid to gas
• States of matter
• Closed system dynamic equilibrium

Safety

• Safety glasses MUST be worn at all times.
• Wear safety gloves when handling iodine.
• Minimise exposure to iodine fumes.

Requirements

2 L measuring cylinder, watch glass, iodine crystals.

Preparation

No prior preparation required.

Demonstration instructions

1. Place some crystals into the cylinder so that they cover the base.
2. Quickly place the watch glass on top of the cylinder to enclose the gas.

Observations

While the change starts immediately, it will take some time to see the purple gas inside the cylinder; it will take up to about 30 min to see a considerable amount of iodine gas. Initially a pale purple gas will be seen in the cylinder. This is most obvious in a well-lit room with a white background.

Chemical equation

2I(s) → I₂(g)

Learning outcomes

• Students should be able to observe that this is an enclosed system where dynamic equilibrium has been set up between the solid and gas states.
• Students should be able to understand that at atmospheric pressure iodine will directly change state from solid to gas.
• Students should be able to write a balanced equation for this process.
• Students should be able to recognise that this is a physical change.

Disposal of chemicals

Place leftover solid iodine back into the iodine bottle.

TOPIC: REACTIONS OF ACIDS AND BASES

Demonstration: Neutralisation reaction

Concepts

• Chemical properties of acids and bases
• pH indicators
• Acid-base reactions producing carbon dioxide gas

Safety

• Safety glasses must be worn at all times.
• Handle all chemicals with care.

Requirements

• 1 L measuring cylinder
• 100 mL water
• 2 large spoonfuls of laundry detergent
• 4 large spoonfuls of NaHCO₃
• 400 mL beaker
• 200 mL vinegar
• 3 mL bromophenol blue indicator
• A large tray to catch any spills

Preparation

In the 1 L measuring cylinder add the water, detergent and NaHCO₃, and in the beaker mix the vinegar with the indicator.

Demonstration instructions

Quickly pour the contents of the beaker into the cylinder.

Observations

• A thick foam is produced.
• When the vinegar is added, the solution in the cylinder changes colour to blue/purple.

Chemical equation

NaHCO₃(s) + CH₃COOH(aq) → CH₃COONa(aq) + H₂O(l) + CO₂(g)

Learning outcomes

• Students should be able to predict the gas that is evolved when an acid reacts with bicarbonate.
• Students should be able to explain the reason for the colour change.
• Students should be able to write a balanced chemical equation for this reaction.

Disposal of chemicals

Dispose of all waste down a laboratory sink with plenty of water.

TOPIC: REDOX REACTIONS

Demonstration: Zinc and copper sulphate redox reaction

Concepts

• Spontaneous redox reactions
• Oxidation and reduction processes
• Balancing redox reactions

Safety

• Safety glasses MUST be worn at all times.
• Handle all chemicals with care.

Equipment

• 2 square zinc strips
• 2 × 200 mL beakers
• 150 mL of 0.1 mol L⁻¹ CuSO₄
• 150 mL of 0.1 mol L⁻¹ ZnSO₄
• Tongs

Preparation

Cut out square zinc strips.

Demonstration instructions

1. Place a zinc strip into each beaker.
2. Pour CuSO₄ solution into one beaker and ZnSO₄ solution into the other beaker.

Observations

• The Zn metal will react with the copper solution immediately and a brown precipitate will form on top of the zinc metal; with time the blue solution of copper sulphate will change to colourless.
• Nothing happens in the second beaker, where zinc is sitting in zinc sulphate solution.
• After 30 minutes more brown precipitate deposits on the zinc in the copper sulphate solution and there is still no reaction in the second beaker.

Chemical equation

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) E° = +1.1 V

Zn(s) + ZnSO₄(aq) → No reaction

Learning outcomes

• Students should understand and be able to explain the reaction between Zn metal and CuSO₄.
• Students should understand and be able to explain why there is no reaction between Zn metal and ZnSO₄.
• Students should be able to write full equations for the redox reactions.
• Students should be able to predict if a reaction is spontaneous or not.
• Students should be able to calculate the potential of a redox reaction.

Disposal of chemicals

Dispose of liquid waste into a laboratory sink with plenty of water. Dispose of solid waste into a laboratory bin.

TOPIC: CHEMICAL PROPERTIES OF ALCOHOLS

Demonstration: Oxidation of alcohol using dichromate

Concepts

Oxidation of alcohols

Safety

• Wash hands immediately if they come into contact with dichromate.
• Keep flames away from alcohol.

Equipment

• Overhead projector
• 2 crystallisation dishes
• 3 mL spirits
• 6 mL acidified 0.1 mol L⁻¹ K₂Cr₂O₇

Demonstration instructions

1. Place both crystallisation dishes on the overhead projector.
2. Pour half of the K₂Cr₂O₇ into each dish.
3. To one dish add 3 mL of the vodka and into the other dish add 3 mL of water.

Observations

After a couple of minutes the dish with the added vodka changes colour from orange to green.

Chemical equation

3CH₃CH₂OH + 16H⁺ + 2Cr₂O₇²⁻ → 3CH₃COOH + 4Cr³⁺ + 11H₂O

(orange Cr₂O₇²⁻ becomes green Cr³⁺)

Learning outcomes

• Students should be able to calculate oxidation numbers for an organic reaction.
• Students should be able to understand the reaction of ethanol with dichromate.
• Students should be able to write a balanced chemical equation to represent this reaction.
• Students should be able to predict the organic product when any alcohol is oxidised.
• Students should be able to explain how breathalysers work.

Disposal of chemicals

Dispose of the dichromate solution into a special waste beaker and follow the school’s protocols for disposal of this chemical. Wash all other chemicals down the laboratory sink with plenty of water.

TOPIC: REDUCING SUGARS

Demonstration: Silver mirror test

Concepts

• Reducing sugars
• Reactions of monosaccharides, disaccharides and polysaccharides
• Reactions of aldehydes

Safety

DO NOT breathe in concentrated ammonia fumes.

Requirements

• 150 mL of 0.1 M AgNO₃
• 2–4 mL of conc. ammonia solution
• 4 g of glucose
• 250 mL round-bottom flask, stopper and pipette

Preparation

No prior preparation.

Demonstration instructions

Place 150 mL of 0.1 M AgNO₃ into a 250 mL round-bottom flask, add 2–4 mL of conc. ammonia solution and then add 4 g of glucose. Stopper the round-bottom flask and swirl the solution for about 30 min. You can speed up the formation of the mirror by placing the flask in a hot water bath (70 °C).

Observations

As the solution is swirled around the flask, the inside of the flask is covered with a reflective material, which looks like a mirror.

Learning outcomes

Students should be able to explain what a reducing sugar is.

Disposal of chemicals

Place any leftover silver nitrate into a special silver collection bottle for correct disposal.

TOPIC: AMINO ACIDS

Demonstration: Iodine fuming (fingerprinting)

Concepts

Amino acids

Safety

DO NOT breathe in the fumes of iodine.

Requirements

• Iodine crystals
• 1 L beaker or chromatography development tank
• Clean A4 sheet of paper
• Vaseline

Preparation

Place some solid iodine at the bottom of a chromatography developing plastic tank (cylinder-like) and cover it; allow the iodine fumes to develop.

Demonstration instructions

Get a clean sheet of white paper, ask a student to place his/her palm on the sheet of paper and then place the paper inside the developing tank (a sweaty palm gives the best results). The finger and palm print develops in a few minutes.

Observations

After a few minutes the imprint of the hand appears as a dark brown stain on the paper. Iodine fuming is based on electron donor–acceptor complexes; basically, a complex is formed between iodine and the amino acids found in the sweat.

Learning outcomes

Students should be able to understand that sweat contains amino acids.

Disposal of chemicals

Place any leftover iodine back into the iodine bottle.

5. Add few drops of phenolphthalein in the content.