Chapter 3 - Metals and Non-Metals Class 10 (Science Solutions)
Ultimate NCERT Solutions for Chapter 3 Metals and Non-Metals Class 10 Science
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NCERT Solutions for Chapter 3 Metals and Non-Metals Class 10 Science
(Question/Answers, Activity & Projects)
Chapter 3 Metals and Non-Metals Class 10 Science Q & A
Activity 3.1
Q 1. Take samples of iron, copper, aluminum and magnesium. Note the appearance of each sample.
Ans 1: Initial Appearance
- Iron: Dull, possibly with a layer of rust if exposed to moisture, appearing reddish-brown.
- Copper: Often tarnished or with a greenish patina if exposed to air, otherwise it has a reddish-brown color.
- Aluminum: Typically has a dull silver appearance with a layer of oxidation, which can make it appear slightly darker.
- Magnesium: Usually appears dull gray due to an oxide layer on the surface.
Q 2. Clean the surface of each sample by rubbing them with sand paper and note their appearance again.
Ans 2: After Cleaning with Sandpaper
- Iron: Becomes shiny and metallic silver in appearance after the rust or oxidation layer is removed.
- Copper: Shows a bright, reddish-pink luster after the tarnish or patina is removed.
- Aluminum: Reveals a bright silver appearance, smoother and more reflective than before.
- Magnesium: Appears shiny and silvery-white once the dull oxide layer is removed.
Cleaning with sandpaper removes the oxide layer or tarnish on each metal, revealing the pure, shiny surface underneath. This process is commonly used to prepare metals for experiments or further testing.
Activity 3.2
Q 1. Take small pieces of iron, copper, aluminum, and magnesium. Try to cut these metals with a sharp knife and note your observations.
Ans 1: When trying to cut small pieces of iron, copper, aluminum, and magnesium with a sharp knife:
- Iron: Very hard and difficult to cut with a knife; usually requires a saw or heavy tool due to its strong metallic bonds.
- Copper: Hard but slightly softer than iron; difficult to cut with a knife and often requires specialized tools.
- Aluminum: Relatively softer than iron and copper, though still challenging to cut with a knife. It might scratch or dent slightly but does not cut easily.
- Magnesium: Softer than the other three metals, slightly easier to cut but still quite challenging to slice with a standard knife.
Q 2. Hold a piece of sodium metal with a pair of tongs. CAUTION: Always handle sodium metal with care. Dry it by pressing between the folds of a filter paper. Put it on a watch-glass and try to cut it with a knife.
Ans 2:
- Using tongs and handling sodium metal with caution, pressing it between filter paper to dry it.
- Placing the sodium on a watch-glass and attempting to cut it with a knife.
Q 3. What do you observe?
Ans 3: Observations on Sodium:
- Sodium is much softer than the other metals mentioned. It can be easily cut with a knife as it has a waxy texture.
- Freshly cut sodium appears shiny, but it quickly dulls due to oxidation when exposed to air.
Activity 3.3
- Take pieces of iron, zinc, lead and copper.
- Place any one metal on a block of iron and strike it four or five times with a hammer. What do you observe?
- Repeat with other metals.
- Record the change in the shape of these metals.
Ans: In Activity 3.3, when you place pieces of iron, zinc, lead, and copper on a block of iron and strike them a few times with a hammer, here’s what you are likely to observe:
- Iron: It is relatively hard and will not deform easily when struck with a hammer. You may observe minor denting, but it retains its original shape better than softer metals.
- Zinc: Zinc is a brittle metal, so striking it may cause it to crack or break instead of flattening or deforming smoothly.
- Lead: Lead is a very soft and malleable metal, so it will flatten and change shape quite easily when struck with a hammer.
- Copper: Copper is also a malleable metal, though it is harder than lead. When struck, it will flatten and spread out somewhat, changing shape, though not as dramatically as lead.
Observe to change in the shape of these metals:
- Iron: Minimal or no change in shape.
- Zinc: Likely to crack or break instead of deforming.
- Lead: Significant flattening and deformation.
- Copper: Noticeable flattening and some spreading out.
This activity demonstrates the varying malleability of metals, with lead and copper being more malleable than iron and zinc.
Activity 3.4
Q 1. List the metals whose wires you have seen in daily life.
Ans 1: Here’s a list of metals commonly seen in wire form in daily life:
- Copper – Used in electrical wiring due to its high conductivity.
- Aluminum – Often used in power transmission lines as it’s lightweight and conductive.
- Iron/Steel – Found in fencing, binding wires, and structural support wires.
- Gold – Used in fine electronic components and high-end cables due to its conductivity and resistance to corrosion.
- Silver – Seen in specialty electronic components due to its high conductivity.
- Nickel – Used in some battery and electronic connections for durability.
- Brass – Sometimes used in decorative wires and musical instruments.
These metals are selected based on their properties like conductivity, durability, and resistance to corrosion.
Activity 3.5
- Take an aluminum or copper wire. Clamp this wire on a stand, as shown in Fig. 3.1.
- Fix a pin to the free end of the wire using wax.
- Heat the wire with a spirit lamp, candle or a burner near the place where it is clamped.
Figure 3.1 Metals are good conductors of heat.
Q 1. What do you observe after some time?
Ans 1: As you heat the wire near the clamped end, after some time, the wax holding the pin begins to melt. Eventually, the pin falls off.
Q 2. Note your observations. Does the metal wire melt?
Ans 2: This is because metals like aluminum or copper are excellent conductors of heat. When you apply heat to one end of the wire, the heat quickly travels through the wire to the free end, where the wax is holding the pin. As the wax heats up, it melts, causing the pin to drop.
No, the wire itself does not melt. This is because metals like aluminum and copper have much higher melting points than wax.
Activity 3.6
- Set up an electric circuit as shown in Fig. 3.2.
- Place the metal to be tested in the circuit between terminals A and B as shown.
Figure 3.2 Metals are good conductors of electricity.
Q 1. Does the bulb glow? What does this indicate?
Ans 1: Observation: When you connect the metal between terminals A and B, check if the bulb in the circuit glows.
Interpretation:
- If the bulb glows, it indicates that the metal conducts electricity, meaning it is a good conductor.
- If the bulb does not glow, it means the metal does not conduct electricity well, suggesting it is a poor conductor or possibly an insulator.
This experiment is a simple way to test the conductivity of various materials.
Q 2. Do metals also conduct electricity? Let us find out
Ans 2: yes, metals are excellent conductors of electricity! This is because metals have free electrons that can move easily through the material. When an electric field is applied, these free electrons flow, creating an electric current. Some of the best conducting metals are silver, copper, and gold.
Here’s a simple way to see this in action:
- Set up a circuit with a battery, wires, and a small bulb or an LED.
- Include the metal in the circuit, using it to connect the bulb to the battery.
- Watch the bulb—if it lights up, the metal is conducting electricity.
This experiment works because metals have a “sea of electrons,” meaning their atoms are surrounded by electrons that aren’t bound tightly to any particular atom, allowing electricity to flow through them easily.
Activity 3.7
- Collect samples of carbon (coal or graphite), Sulphur and iodine.
- Carry out the Activities 3.1 to 3.4 and 3.6 with these non-metals and record your observations.
- Compile your observations regarding metals and non-metals in Table 3.1.
Ans: Here’s a table format you can use for recording observations on the non-metals carbon (coal or graphite), sulfur, and iodine during Activities 3.1 to 3.4 and 3.6:
Element | Symbol | Type of surface | Hardness | Malleability | Ductility | Conducts Electricity | Sonority (Sound) |
Carbon (Coal/Graphite) | C | rough | Varies (Graphite: Soft, Coal: Hard) | Graphite: Brittle, Coal: Brittle | Not ductile | Graphite conducts, coal does not | Low sonority (Coal), Moderate (Graphite) |
Sulfur | S | Rough | Soft | Brittle | Not ductile | Does not conduct | Low sonority |
Iodine | I | Shiny | Brittle | Brittle | Not ductile | Does not conduct | Low sonority |
- Hardness: Describes how hard the material is. Graphite is softer, coal may vary, sulfur and iodine are relatively soft and brittle.
- Malleability: Whether the element can be hammered into thin sheets. Non-metals like sulfur and iodine are brittle and not malleable.
- Ductility: Whether the element can be stretched into wires. Non-metals like sulfur and iodine are not ductile.
- Conducts Electricity: Graphite conducts electricity due to its free electrons, whereas sulfur and iodine do not conduct electricity.
- Sonority: Refers to the sound produced when struck. Non-metals like sulfur and iodine produce a low sound (dull), while graphite has a moderate sound.
Activity 3.8 [Experiment]
- Take a magnesium ribbon and some Sulphur powder.
- Burn the magnesium ribbon. Collect the ashes formed and dissolve them in water.
- Test the resultant solution with both red and blue litmus paper.
- Is the product formed on burning magnesium acidic or basic?
- Now burn Sulphur powder. Place a test tube over the burning Sulphur to collect the fumes produced.
- Add some water to the above test tube and shake.
- Test this solution with blue and red litmus paper.
- Is the product formed on burning Sulphur acidic or basic?
- Can you write equations for these reactions?
Solution: Let’s go step by step through each part of your experiment and explain the reactions:
- Burning Magnesium Ribbon:
- When magnesium (Mg) is burned in air, it reacts with oxygen (O₂) to form magnesium oxide (MgO).
- Reaction:
2Mg(s) + O₂(g) → 2MgO(s) - Magnesium oxide is a basic oxide. When you dissolve the ashes of the magnesium ribbon (which is magnesium oxide) in water, it forms magnesium hydroxide, which is a base.
- Reaction with water:
MgO(s) + H₂O(l) → Mg(OH)₂(aq) - The resulting solution will turn red litmus paper blue, indicating that the solution is basic.
- Burning Sulphur Powder:
- When sulfur (S) burns, it reacts with oxygen to form sulfur dioxide (SO₂).
- Reaction:
S(s) + O₂(g) → SO₂(g) - You collect the fumes (SO₂) in a test tube. When you add water to it, sulfur dioxide dissolves in the water and reacts to form sulfurous acid (H₂SO₃).
- Reaction with water:
SO₂(g) + H₂O(l) → H₂SO₃(aq) - Sulfurous acid is an acidic solution, so when you test it with blue litmus paper, it will turn red, indicating that the solution is acidic.
Conclusion:
- The product formed on burning magnesium is basic (magnesium hydroxide).
- The product formed on burning sulfur is acidic (sulfurous acid).
Summary of Equations:
- Magnesium Burning:
2Mg(s) + O₂(g) → 2MgO(s)
MgO(s) + H₂O(l) → Mg(OH)₂(aq)
- Sulfur Burning:
S(s) + O₂(g) → SO₂(g)
SO₂ (g) + H₂O (l) → H₂SO₃ (aq)
QUESTIONS
Q 1. Give an example of a metal which:
- Is a liquid at room temperature;
- Can be easily cut with a knife;
- Is the best conductor of heat;
- Is a poor conductor of heat.
Ans 1:
1. Liquid at room temperature:
- Mercury (Hg) is the only metal that is a liquid at room temperature (around 20°C or 68°F).
2. Can be easily cut with a knife:
- Sodium (Na) is a soft alkali metal that can be easily cut with a knife.
3. Best conductor of heat:
- Silver (Ag) is the best conductor of heat among all metals.
4. Poor conductor of heat:
- Titanium (Ti) is a metal that is a relatively poor conductor of heat compared to other metals like copper or silver.
Q 2. Explain the meanings of malleable and ductile.
Ans 2: Malleable refers to a material’s ability to be hammered or pressed into thin sheets without breaking. It indicates flexibility when subjected to compressive force.
Ductile refers to a material’s ability to stretch into a wire when pulled. It shows flexibility under tensile stress.
Activity 3.9 [ Do Experiment]
- CAUTION: The following activity needs the teacher’s assistance. It would be better if students wear eye protection.
- Hold any of the samples taken above with a pair of tongs and try burning over a flame. Repeat with the other metal samples.
- Collect the product if formed.
- Let the products and the metal surface cool down.
- Which metals burn easily?
- What flame colour did you observe when the metal burnt?
- How does the metal surface appear after burning?
- Arrange the metals in the decreasing order of their reactivity towards oxygen.
- Are the products soluble in water?
Solution 3.9: This is a great experiment to explore the reactivity of different metals with oxygen, which helps students understand oxidation and metal behavior in different environments. Here’s how you could guide students through the activity:
Activity Instructions
- Safety First: Ensure that students wear eye protection (safety goggles) and handle the metals with care, especially near the flame. It’s advisable that a teacher assists throughout the process.
- Heating Metals: Using tongs, hold a small piece of each metal sample (e.g., magnesium, zinc, copper, etc.) and carefully heat them in a flame (from a Bunsen burner or a similar controlled heat source). Observe the reaction when the metals burn.
- Product Collection: After the metal burns, collect the resulting product, which may be a metal oxide or another compound formed due to the reaction with oxygen in the air. Allow the products and metal pieces to cool down before further handling.
- Observations:
- Ease of Burning: Which metal burns easily? Generally, metals like magnesium burn easily with a bright white flame, while others like copper may burn with a green flame.
- Flame Colour: Record the color of the flame observed when each metal burns (e.g., magnesium burns with a bright white light, copper burns with a blue-green flame).
- Appearance After Burning: Look at the appearance of the metal surface after burning. Some metals, like magnesium, turn white (as magnesium oxide), while others like iron may show signs of rusting or a dull coating after burning.
5. Reactivity Ranking:
- After observing the burning behavior, arrange the metals in the decreasing order of their reactivity towards oxygen. Metals like magnesium and zinc are more reactive than copper, which reacts more slowly in comparison.
6. Solubility in Water: Test whether the metal oxides (or other products) are soluble in water. For example:
- Magnesium oxide dissolves in water to form magnesium hydroxide.
- Zinc oxide may show slight solubility in water.
- Copper oxide is generally insoluble in water.
Key Concepts to Discuss:
- Reactivity Series: This experiment gives a visual representation of the reactivity of metals. More reactive metals tend to burn more easily and form metal oxides quickly.
- Oxidation: The reaction with oxygen (combustion) produces metal oxides, illustrating the concept of oxidation.
- Flame Colors: The different flame colors are due to the different energy levels and the wavelengths of light emitted by each metal when heated.
After completing the activity, students can discuss the results and understand how metals vary in their reactivity and how this relates to their placement in the reactivity series.
Activity 3.10 [Do Experiment]
CAUTION: This Activity needs the teacher’s assistance.
- Collect the samples of the same metals as in Activity 3.9.
- Put small pieces of the samples separately in beakers half-filled with cold water.
- Which metals reacted with cold water? Arrange them in the increasing order of their reactivity with cold water.
- Did any metal produce fire on water?
- Does any metal start float after some time?
- Put the metals that did not react with cold water in beakers half-filled with hot water.
- For the metals that did not react with hot water, arrange the apparatus as shown in Fig. 3.3 and observe their reaction with steam.
- Which metals did not react even with steam?
- Arrange the metals in the decreasing order of reactivity with water.
Solution: Activity 3.10 [ Student do yourself with the help of Teacher]
[Hint solution]
To investigate the reactivity of metals with water in various forms (cold water, hot water, and steam), follow the steps below:
Step-by-Step Procedure
1. Collect Metal Samples: Use small pieces of metals like sodium, calcium, magnesium, aluminum, iron, zinc, and copper, as referenced in Activity 3.9.
2. Cold Water Test:
- Place each metal sample in a separate beaker half-filled with cold water.
- Observe which metals react immediately, indicating higher reactivity.
3. Observations in Cold Water:
- Metals Reacting with Cold Water: Sodium and calcium generally react with cold water. Sodium reacts vigorously, often producing heat, and may even catch fire on the water’s surface.
- Floating Behavior: Sodium may float due to the hydrogen gas released, which reduces its density.
4. Hot Water Test:
- For metals that did not react with cold water (like magnesium), place them in beakers half-filled with hot water and observe any reactions.
5. Steam Test:
- For metals that did not react with hot water (such as aluminum, zinc, and iron), set up an apparatus to expose the metal samples to steam (as shown in Fig. 3.3).
- Observe which metals react with steam.
6. Final Observations:
- Metals like copper generally do not react with steam or water at any temperature, indicating low reactivity.
Results and Conclusions
- Increasing Order of Reactivity with Cold Water:
- Sodium > Calcium > Magnesium (doesn’t react significantly with cold water but does with hot water).
- Fire Production: Sodium may catch fire when reacting with cold water.
- Floating Behavior: Sodium may float on water due to hydrogen gas production.
- Reaction with Steam:
- Magnesium reacts with hot water and steam.
- Aluminum, zinc, and iron may react with steam but not with cold or hot water.
- Non-Reacting Metals:
- Copper and other less reactive metals do not react with steam.
- Decreasing Order of Reactivity with Water:
- Sodium > Calcium > Magnesium > Aluminum > Zinc > Iron > Copper
This activity demonstrates the reactivity series of metals with water, showing how some metals react with cold water, others with hot water, and some only with steam, while a few do not react at all.
Activity 3.11
- Collect all the metal samples except sodium and potassium again. If the samples are tarnished, rub them clean with sand paper. CAUTION: Do not take sodium and potassium as they react vigorously even with cold water.
- Put the samples separately in test tubes containing dilute hydrochloric acid.
- Suspend thermometers in the test tubes, so that their bulbs are dipped in the acid.
- Observe the rate of formation of bubbles carefully.
- Which metals reacted vigorously with dilute hydrochloric acid?
- With which metal did you record the highest temperature?
- Arrange the metals in the decreasing order of reactivity with dilute acids.
Solution: Activity 3.11
Here’s a step-by-step approach to conducting this activity to observe the reactivity of metals with dilute hydrochloric acid:
Materials Needed
- Metal samples (e.g., magnesium, aluminum, zinc, iron, copper).
- Dilute hydrochloric acid (HCl).
- Sandpaper.
- Test tubes.
- Thermometers.
- Safety goggles and gloves.
Procedure
- Prepare the Metal Samples:
- Collect samples of the metals except for sodium and potassium, as they react too vigorously with water and dilute acids.
- If the metal samples are tarnished (have an oxidized layer), rub them with sandpaper to clean their surfaces and expose the pure metal.
- Set Up the Experiment:
- Place each metal sample separately in a test tube.
- Add dilute hydrochloric acid to each test tube, ensuring enough acid to cover the metal.
- Suspend a thermometer in each test tube so that the bulb is immersed in the acid but not touching the metal.
- Observe the Reactions:
- Watch the formation of bubbles on the metal surfaces, which indicates the production of hydrogen gas. The faster the bubbles form, the more vigorous the reaction.
- Record the temperature increase for each metal. The more reactive the metal, the greater the temperature rise should be.
- Record Observations:
- Identify which metals reacted the most vigorously by observing the intensity of bubbling and the temperature change.
Analysis and Conclusion
- Identify Vigorous Reactions: Metals like magnesium and zinc are expected to react more vigorously with dilute hydrochloric acid, producing more bubbles and increasing the temperature more rapidly than less reactive metals like iron and copper.
- Highest Temperature: The metal that reacts most exothermically (such as magnesium) will likely show the highest temperature increase.
- Reactivity Order:
- Based on your observations, arrange the metals in decreasing order of reactivity. A typical order from highest to lowest reactivity with dilute HCl might be: Magnesium > Zinc > Iron > Copper
Activity 3.12
- Take a clean wire of copper and an iron nail.
- Put the copper wire in a solution of iron sulphate and the iron nail in a solution of copper sulphate taken in test tubes(Fig. 3.4)
- Record your observations after 20 minutes.
- In which test tube did you find that a reaction has occurred?
- On what basis can you say that a reaction has actually taken place?
- Can you correlate your observations for the Activities 3.9, 3.10 and 3.11?
- Write a balanced chemical equation for the reaction that has taken place.
- Name the type of reaction.
Figure 3.4 Reaction of metals with salt solutions
Solution: Activity 3.12
Steps:
- Take a clean copper wire and an iron nail.
- Put the copper wire in a solution of iron sulphate (FeSO₄) and the iron nail in a solution of copper sulphate (CuSO₄) in separate test tubes (refer to Fig. 3.4 in your textbook).
- Wait for 20 minutes and observe.
Observations: After 20 minutes, you may notice:
- Test tube with copper wire in iron sulphate solution: No visible change occurs, meaning no reaction has taken place.
- Test tube with iron nail in copper sulphate solution: A reddish-brown layer forms on the iron nail. This indicates that a reaction has occurred.
Analysis:
- Reaction Occurrence: A reaction has occurred in the test tube where the iron nail was placed in the copper sulphate solution. The reddish-brown layer on the iron nail is copper metal, indicating that iron has displaced copper from copper sulphate.
- Basis for Concluding a Reaction: The change in color and formation of a new substance (copper) on the iron nail confirms that a chemical reaction has taken place.
Correlation with Activities 3.9, 3.10, and 3.11: These activities demonstrate displacement reactions, where a more reactive metal displaces a less reactive metal from its compound in solution.
Chemical Equation: The reaction in the test tube with the iron nail and copper sulphate solution is:
Fe (s) + CuSO₄ (aq) → FeSO₄ (aq) + Cu (s)
Type of Reaction: This is a displacement reaction, where iron displaces copper from its compound due to iron’s higher reactivity compared to copper.
QUESTIONS
Q1. Why is sodium kept immersed in kerosene oil?
Ans 1:
- Sodium is a highly reactive metal, especially with oxygen and moisture in the air.
- Exposure to air can cause sodium to react rapidly, forming sodium oxide and even catching fire.
- Kerosene oil acts as a barrier, preventing sodium from reacting with air or moisture.
- Immersing sodium in kerosene ensures safe storage and prevents accidental ignition or explosions.
Q 2. Write equations for the reactions of
(i) Iron with steam
(ii) calcium with water
(iii) potassium with water
Ans 2: Here are the chemical equations for the reactions:
- Iron with steam: 3Fe + 4H₂O → Fe₃O₄ + 4H₂
- Calcium with water: Ca + 2H₂O → Ca(OH)₂ + H₂
- Potassium with water: 2K + 2H₂O → 2KOH + H₂
Q 3. Samples of four metals A, B, C and D were taken and added to the following solution one by one. The results obtained have been tabulated as follows.
Metal | Iron(II) sulphate | Copper(II) sulphate | Zinc sulphate | Silver nitrate |
A | No reaction | Displacement | ||
B | Displacement | |||
C | No reaction | No reaction | No reaction | Displacement |
D | No reaction | No reaction | No reaction | No reaction |
Use the Table above to answer the following questions about metals A, B, C and D.
(i) Which is the most reactive metal?
Ans (i): The most reactive metal will be the one that displaces the most ions from the solutions. Based on the table:
- Metal B displaces ions from Iron(II) sulphate, suggesting that it is more reactive than iron.
- Metal A displaces ions from Copper(II) sulphate, indicating it is more reactive than copper.
- Metal C displaces ions from silver nitrate, which indicates it is more reactive than silver.
Thus, Metal B is the most reactive metal because it displaces ions from Iron(II) sulphate, which is a less reactive metal.
(ii) What would you observe if B is added to a solution of Copper (II) sulphate?
Ans (ii): Metal B displaces iron from Iron(II) sulphate, so it is more reactive than iron. However, since metal B has no reaction with Copper(II) sulphate, there will be no observable reaction when B is added to Copper(II) sulphate.
(iii) Arrange the metals A, B, C and D in the order of decreasing reactivity
Ans (iii): To determine the reactivity order, we look at the displacement reactions and compare which metals can displace others:
- Metal B can displace iron from Iron(II) sulphate, so it is more reactive than iron. It is the most reactive metal.
- Metal A can displace copper from Copper(II) sulphate, so it is more reactive than copper but less reactive than Metal B.
- Metal C can displace silver from silver nitrate, so it is more reactive than silver but less reactive than both metals A and B.
- Metal D shows no displacement reactions, indicating it is the least reactive.
So, the order of decreasing reactivity is: B > A > C > D
Q 4. Which gas is produced when dilute hydrochloric acid is added to a reactive metal? Write the chemical reaction when iron reacts with dilute H₂SO₄
Ans 4: When dilute hydrochloric acid (HCl) is added to a reactive metal, hydrogen gas (H₂) is produced. This is a typical reaction between an acid and a metal, where the metal displaces hydrogen from the acid.
The general reaction is: Metal + Hydrochloric acid → Salt + Hydrogen gas
For example, when iron (Fe) reacts with dilute sulfuric acid (H₂SO₄), the reaction can be written as:
Fe (s) + H₂SO₄ (aq) → FeSO₄ (aq) + H₂ (g)
Here, iron (Fe) reacts with sulfuric acid (H₂SO₄) to form iron(II) sulfate (FeSO₄) and hydrogen gas (H₂).
Q 5. What would you observe when zinc is added to a solution of iron(II) sulphate? Write the chemical reaction that takes place
Ans 5: When zinc is added to a solution of iron(II) sulfate (FeSO₄), a displacement reaction takes place. Zinc is more reactive than iron and displaces iron from its compound, forming zinc sulfate and iron.
The chemical reaction is: Zn(s) + FeSO₄(aq) → ZnSO₄(aq) + Fe(s)
Observations:
- Iron (Fe) is displaced and appears as a solid, so the solution will have iron deposited on the surface of the zinc.
- The color of the solution will change. Iron(II) sulfate is typically pale green, and as zinc displaces the iron, the solution may become colorless or show a pale-yellow tint due to the formation of zinc sulfate.
Activity 3.13
Q 1. What is the physical state of these salts?
Ans 1: Physical State of Salts:
- Sodium chloride (NaCl): Solid, crystalline.
- Potassium iodide (KI): Solid, crystalline.
- Barium chloride (BaCl₂): Solid, crystalline
Q 2. Take a small amount of a sample on a metal spatula and heat directly on the flame (Fig. 3.7). Repeat with other samples.
Ans 2: When heating these salts directly on a flame, you may observe the following:
- Sodium chloride: The flame does not impart a distinct color, as NaCl is colorless in the flame.
- Potassium iodide: Potassium salts impart a lilac/purple color to the flame due to the excitation of potassium ions.
- Barium chloride: Barium salts impart a green color to the flame due to the excitation of barium ions.
Figure 3.7 Heating a salt sample on a spatula
Q 3. What did you observe? Did the samples impart any colour to the flame? Do these compounds melt?
Ans 3: These salts do not typically melt on direct heating with a metal spatula under a standard laboratory flame. They require much higher temperatures (much above the typical flame) to melt, indicating they have high melting points.
Chapter 3 – Metals and Non-metals Class 10 Science Solutions Pdf Download
Updated Solution 2024-2025
Q 4. Try to dissolve the samples in water, petrol and kerosene. Are they soluble?
Ans 4: Solubility in Different Solvents:
- Water:
- Sodium chloride: Soluble in water, forms a clear solution.
- Potassium iodide: Soluble in water, forms a clear solution.
- Barium chloride: Soluble in water, forms a clear solution.
- Petrol:
- These salts are generally insoluble in petrol due to the non-polar nature of petrol and the ionic nature of the salts.
- Kerosene:
- Similar to petrol, these salts are insoluble in kerosene for the same reason.
Q 5. Make a circuit as shown in Fig. 3.8 and insert the electrodes into a solution of one salt. What did you observe? Test the other salt samples too in this manner.
Ans 5: When making a circuit and inserting electrodes into a solution of these salts (e.g., dissolved in water), you will observe:
- Sodium chloride: The solution conducts electricity, as NaCl dissociates into Na⁺ and Cl⁻ ions, which are free to move and carry the current.
- Potassium iodide: Similar to NaCl, KI dissociates into K⁺ and I⁻ ions in water, allowing the solution to conduct electricity.
- Barium chloride: Barium chloride dissociates into Ba²⁺ and Cl⁻ ions, enabling the solution to conduct electricity.
Figure 3.8 Testing the conductivity of a salt solution
Q 6. What is your inference about the nature of these compounds?
Ans 6: The observed properties suggest that these salts are ionic compounds. This is evident from their:
- High melting points.
- Solubility in water (a polar solvent).
- Ability to conduct electricity when dissolved in water (electrolyte behavior).
- Distinct flame colors indicating the presence of specific metal ions (Na⁺, K⁺, Ba²⁺).
These characteristics point to the ionic nature of these compounds, where ions are held together by strong ionic bonds, making them typically soluble in water and conductive in solution.
Questions
Q 1. (i) write the electron-dot structures for sodium, oxygen and magnesium.
Ans (i): Here are the electron-dot (Lewis’s dot) structures for Sodium (Na), Oxygen (O), and Magnesium (Mg):
1. Sodium (Na):
- Sodium has an atomic number of 11, so its electron configuration is 1s² 2s² 2p⁶ 3s¹.
- The valence electron is in the 3rd shell (3s¹).
- The electron-dot structure for sodium is:
2. Oxygen (O):
- Oxygen has an atomic number of 8, so its electron configuration is 1s² 2s² 2p⁴.
- The valence electrons are in the 2nd shell (2s² 2p⁴), and oxygen has 6 valence electrons.
- The electron-dot structure for oxygen is:
Here, the two dots on each side represent the lone pairs, and the two remaining dots form a bond if oxygen is involved in a molecule (e.g., H₂O).
3. Magnesium (Mg):
- Magnesium has an atomic number of 12, so its electron configuration is 1s² 2s² 2p⁶ 3s².
- The valence electrons are in the 3rd shell (3s²), so magnesium has 2 valence electrons.
- The electron-dot structure for magnesium is:
These dot structures represent the valence electrons of each element.
Q (ii). Show the formation of Na₂O and MgO by the transfer of electrons.
Ans (ii): The formation of Na₂O (sodium oxide) and MgO (magnesium oxide) involves the transfer of electrons to form ionic bonds:
- Na₂O (Sodium Oxide):
- Sodium (Na) has one electron in its outer shell. It loses this electron to become Na⁺.
- Oxygen (O) has six electrons in its outer shell and needs two more to complete its octet. It gains two electrons to become O²⁻.
- Two Na atoms each lose one electron, and one O atom gains two electrons:
2Na → 2Na⁺ + 2e⁻
O + 2e⁻ → O²⁻
The result is Na₂O: 2Na⁺ + O²⁻
MgO (Magnesium Oxide):
- Magnesium (Mg) has two electrons in its outer shell. It loses both to become Mg²⁺.
- Oxygen (O) gains two electrons to become O²⁻.
- One Mg atom loses two electrons, and one O atom gains two electrons
Mg → Mg²⁺ + 2e⁻
O + 2e⁻ → O²⁻
The result is MgO: Mg²⁺ + O²⁻
Q (iii). What are the ions present in these compounds?
Ans (iii): Ions present in Na₂O:
- Sodium ions (Na⁺)
- Oxide ions (O²⁻)
Ions present in MgO:
- Magnesium ions (Mg²⁺)
- Oxide ions (O²⁻)
Q 2. Why do ionic compounds have high melting points?
Ans 2: Ionic compounds have high melting points because the strong electrostatic forces between positively and negatively charged ions require a large amount of energy to break.
Questions
Q 1. Define the following terms.
(i) Mineral: A naturally occurring inorganic substance with a definite chemical composition and crystal structure.
(ii) Ore: A type of rock that contains valuable minerals, typically metals, which can be extracted profitably.
(iii) Gangue: The unwanted or worthless material found alongside valuable minerals in an ore.
Q 2. Name two metals which are found in nature in the free state
Ans 2: Two metals found in nature in the free state are gold and platinum.
Q 3. What chemical process is used for obtaining a metal from its oxide?
Ans 3: The chemical process used to obtain a metal from its oxide is called reduction. In this process, the metal oxide is heated with a reducing agent, such as carbon, which removes the oxygen from the metal oxide, leaving the pure metal.
Activity 3.14
- take three test tubes and place clean iron nails in each of them.
- Label these test tubes A, B and C. Pour some water in test tube A and cork it.
- Pour boiled distilled water in test tube B, add about 1 mL of oil and cork it. The oil will float on water and prevent the air from dissolving in the water.
- Put some anhydrous calcium chloride in test tube C and cork it. Anhydrous calcium chloride will absorb the moisture, if any, from the air. Leave these test tubes for a few days and then observe (Fig. 3.13).
Solution: Activity 3.14 [Experiment] (students do yourself with the help of teacher)
{hint Answer}: This experiment seems designed to observe the effects of different conditions on the oxidation of iron nails, based on varying the amount of moisture and air available. Here’s a breakdown of what you can expect from each test tube and how the conditions will impact the iron nails:
1. Test Tube A (Water)
- Condition: This test tube contains water and the nails are submerged in it.
- Expectation: Since water is present, and air can dissolve in water, the iron nails in this test tube will undergo oxidation (rusting) over time. This is a typical condition for rust formation, as both water and oxygen are needed for the process.
2. Test Tube B (Boiled Distilled Water with Oil)
- Condition: This test tube contains boiled distilled water and oil, which floats on top of the water, preventing air from dissolving in the water.
- Expectation: Because the oil prevents the absorption of oxygen from the air into the water, the iron nails in this test tube will likely rust more slowly, or may not rust at all, as oxygen is less available to react with the iron.
3. Test Tube C (Anhydrous Calcium Chloride)
- Condition: This test tube contains anhydrous calcium chloride, which absorbs any moisture from the air.
- Expectation: Since anhydrous calcium chloride absorbs moisture, the test tube’s environment will be dry, and the iron nails will likely not rust, or will rust at a very slow rate, because the moisture required for oxidation will be absent.
Observations after a few days:
- The nails in Test Tube A will show noticeable rusting due to the presence of water and oxygen.
- The nails in Test Tube B may have little to no rusting, as the lack of oxygen in the water limits the oxidation process.
- The nails in Test Tube C will likely show little to no rusting, since the absence of moisture from the air prevents oxidation.
This experiment demonstrates the importance of both moisture and oxygen in the rusting process of iron.
Figure 3.13 Investigating the conditions under which iron rusts. In tube A, both air and water are present. In tube B, there is no air dissolved in the water. In tube C, the air is dry
Questions
Q 1. Metallic oxides of zinc, magnesium and copper were heated with the following metals.
Metal | Zinc | Magnesium | Copper |
Zinc oxide |
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|
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Magnesium oxide |
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Copper oxide |
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In which cases will you find displacement reactions taking place?
Ans 1: Here’s a table showing the expected results when metallic oxides of zinc, magnesium, and copper are heated with the metals zinc, magnesium, and copper. Displacement reactions will occur if the metal added is more reactive than the metal in the oxide, as a more reactive metal can displace a less reactive metal from its compound.
Table: Reaction of Metal Oxides with Different Metals
Metal | Zinc | Magnesium | Copper |
Zinc oxide | No reaction | Displacement | No reaction |
Magnesium oxide | No reaction | No reaction | No reaction |
Copper oxide | Displacement | Displacement | No reaction |
Explanation:
- Zinc Oxide: Magnesium, being more reactive than zinc, can displace zinc from zinc oxide, so a displacement reaction occurs when magnesium is used. No reaction occurs with zinc or copper because they are not more reactive than zinc.
- Magnesium Oxide: No reaction occurs with zinc, magnesium, or copper since none are more reactive than magnesium.
- Copper Oxide: Zinc and magnesium can displace copper from copper oxide because they are more reactive than copper, but copper itself does not react.
In summary, displacement reactions will occur:
- Zinc Oxide with Magnesium
- Copper Oxide with Zinc and Magnesium
Q 2. Which metals do not corrode easily?
Ans 2: Metals that do not corrode easily are typically those that are highly resistant to oxidation and environmental reactions. These include:
- Gold – Known for its excellent resistance to tarnish and corrosion, even in harsh environments. Gold is widely used in jewelry and electronics for this reason.
- Platinum – This metal is extremely resistant to corrosion and is also used in jewelry, as well as in industrial applications, including catalytic converters.
- Titanium – Highly resistant to corrosion from saltwater and chlorine, making it valuable for medical implants and marine applications.
- Silver – Although it tarnishes (a form of corrosion) in the presence of sulfur compounds, it generally resists other forms of corrosion well.
These metals either form a protective oxide layer naturally or are noble metals, meaning they do not react easily with their environment.
Q 3. What are alloys?
Ans 3: Alloys are materials made by combining two or more metals or a metal with a non-metal to enhance properties like strength, durability, or resistance to corrosion. Examples include steel (iron and carbon) and bronze (copper and tin).
EXERCISE
Q 1. Which of the following pairs will give displacement reactions?
(a) NaCl solution and copper metal
(b) MgCl₂ solution and aluminum metal
(c) FeSO₄ solution and silver metal
(d) AgNO₃ solution and copper metal
Ans 1: (d). AgNO₃ solution and copper metal.
Q 2. Which of the following methods is suitable for preventing an iron frying pan from rusting?
- Applying grease
- Applying paint
- Applying a coating of zinc
- All of above
Ans 2: c) Applying a coating of zinc: This process, known as galvanization, is commonly used to protect iron or steel from rusting. While grease and paint can provide some protection, they are not as effective or durable as zinc coating for preventing rust on iron frying pans, especially with frequent use.
Q 3. An element reacts with oxygen to give a compound with a high melting point. This compound is also soluble in water. The element is likely to be:
- Calcium
- Carbon
- Silicon
- Iron
Ans 3: a) Calcium. Calcium reacts with oxygen to form calcium oxide, a compound with a high melting point that is also soluble in water, forming a basic solution (calcium hydroxide).
Q 4. Food cans are coated with tin and not with zinc because
- zinc is costlier than tin.
- zinc has a higher melting point than tin.
- zinc is more reactive than tin.
- zinc is less reactive than tin.
Ans 4: (c) zinc is more reactive than tin: Food cans are coated with tin instead of zinc because zinc is more reactive and could corrode or react with the food inside, which may contaminate it. Tin, being less reactive, provides a safer, stable coating.
Q 5. You are given a hammer, a battery, a bulb, wires and a switch.
a) How could you use them to distinguish between samples of metals and non-metals?
b) Assess the usefulness of these tests in distinguishing between metals and non-metals
Ans 5: a) You can create a simple circuit with the battery, wires, bulb, and switch. Connect the metal sample to the circuit. If the bulb lights up, the material is a metal (good conductor). If it doesn’t, the material is a non-metal (poor conductor).
b) This test is useful for distinguishing between metals (which conduct electricity) and non-metals (which do not), but it may not work for non-metals that conduct under certain conditions (e.g., graphite).
Q 6. What are amphoteric oxides? Give two examples of amphoteric oxides.
Ans 6: Amphoteric oxides are chemical compounds that can react both as acids and as bases, depending on the conditions. This means they can react with both acids and bases to form salts and water, exhibiting dual behavior.
Examples of Amphoteric Oxides:
- Aluminum oxide (Al₂O₃): It reacts with acids, like hydrochloric acid (HCl), to form aluminum chloride (AlCl₃) and water, and with bases, like sodium hydroxide (NaOH), to form sodium aluminate.
- Reaction with acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
- Reaction with base: Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄]
- Zinc oxide (ZnO): Zinc oxide reacts with acids, like sulfuric acid (H₂SO₄), to form zinc sulfate (ZnSO₄) and water, and with bases, like sodium hydroxide (NaOH), to form sodium zincate.
- Reaction with acid: ZnO + H₂SO₄ → ZnSO₄ + H₂O
- Reaction with base: ZnO + 2NaOH + H₂O → Na₂[Zn(OH)₄
Q 7. Name two metals which will displace hydrogen from dilute acids, and two metals which will not.
Ans 7: Two metals that will displace hydrogen from dilute acids are:
- Zinc (Zn)
- Magnesium (Mg)
These metals are more reactive than hydrogen and can displace it from acids such as hydrochloric acid (HCl) or sulfuric acid (H₂SO₄).
Two metals that will not displace hydrogen from dilute acids are:
- Gold (Au)
- Platinum (Pt)
These metals are less reactive than hydrogen and do not easily undergo reactions with dilute acids to displace hydrogen.
Q 8. In the electrolytic refining of a metal M, what would you take as the anode, the cathode and the electrolyte?
Ans 8: In the electrolytic refining of a metal M, the setup involves the following:
- Anode: The impure metal M is taken as the anode. This is because, during electrolysis, the impure metal at the anode undergoes oxidation and dissolves into the electrolyte as metal ions.
- Cathode: The pure metal M is deposited at the cathode. The metal ions from the electrolyte gain electrons (reduction) at the cathode and are deposited as pure metal.
- Electrolyte: The electrolyte typically contains a solution of a salt of metal M, such as its sulfate (e.g., copper sulfate for copper refining) or chloride (e.g., copper chloride for copper refining). The electrolyte provides a medium for the movement of metal ions between the anode and cathode.
Thus, in the electrolytic refining process, the impure metal M is used as the anode, pure metal M is deposited at the cathode, and the electrolyte contains the metal salt of M.
Q 9. Pratyush took Sulphur powder on a spatula and heated it. He collected the gas evolved by inverting a test tube over it, as shown in figure below.
(a) What will be the action of gas on
- (i) dry litmus paper?
- (ii) moist litmus paper?
(b) Write a balanced chemical equation for the reaction taking place.
Ans 9: (a) Action of gas:
- (i) On dry litmus paper: No effect.
- (ii) On moist litmus paper: Turns red (indicating an acidic gas).
(b) Balanced chemical equation: S(s) + O₂(g) → SO₂(g)
Q 10. State two ways to prevent the rusting of iron.
Ans 10:
- Coating with paint or oil: This creates a protective barrier, preventing exposure to moisture and oxygen.
- Galvanization: Coating iron with a layer of zinc prevents rust by protecting the iron from corrosion.
Q 11. What type of oxides are formed when non-metals combine with oxygen?
Ans 11: When non-metals combine with oxygen, they form non-metal oxides, which are typically acidic in nature. These oxides can react with water to form acids. Examples include carbon dioxide (CO₂) and sulfur dioxide (SO₂).
Q 12. Give reasons
- Platinum, gold and silver are used to make jewellery.
- Sodium, potassium and lithium are stored under oil.
- Aluminum is a highly reactive metal, yet it is used to make utensils for cooking.
- Carbonate and sulphide ores are usually converted into oxides during the process of extraction.
Ans 12: (a) Platinum, gold, and silver are used for jewelry because they are rare, durable, and have an attractive appearance.
(b) Sodium, potassium, and lithium are stored under oil to prevent reactions with moisture and air, which could cause them to react violently.
(c) Aluminum is highly reactive, but it forms a protective oxide layer on its surface that prevents further reaction, making it safe and suitable for cooking utensils.
(d) Carbonate and sulphide ores are converted into oxides during extraction to make them easier to reduce into pure metals using heat or chemicals.
Q 13. You must have seen tarnished copper vessels being cleaned with lemon or tamarind juice. Explain why these sour substances are effective in cleaning the vessels.
Ans 13: Lemon and tamarind juice are effective in cleaning tarnished copper vessels because they contain acids, primarily citric acid and tartaric acid. These acids react with the copper oxide (tarnish) on the surface of the vessel, breaking it down and dissolving it. The acid also helps to loosen dirt and stains, making it easier to clean the copper and restore its shine.
Q 14. Differentiate between metal and non-metal on the basis of their chemical properties.
Ans 14: Metals and non-metals differ in chemical properties as follows:
- Reaction with Oxygen: Metals form basic oxides (e.g., Na₂O), while non-metals form acidic oxides (e.g., CO₂).
- Reaction with Water: Metals like sodium react with water to form hydroxides and release hydrogen gas. Non-metals generally do not react with water.
- Reaction with Acids: Metals react with acids to produce hydrogen gas, while non-metals do not typically react with acids.
- Conductivity: Metals are good conductors of electricity and heat, while non-metals are poor conductors.
- Electronegativity: Non-metals have higher electronegativity, attracting electrons more strongly than metals.
Q 15. A man went door to door posing as a goldsmith. He promised to bring back the glitter of old and dull gold ornaments. An unsuspecting lady gave a set of gold bangles to him which he dipped in a particular solution. The bangles sparkled like new but their weight was reduced drastically. The lady was upset but after a futile argument the man beat a hasty retreat. Can you play the detective to find out the nature of the solution he had used?
Ans 15: The solution used by the man in this scenario is likely to be a strong acid, such as aqua regia (a mixture of hydrochloric acid and nitric acid), or a concentrated nitric acid solution. Here’s why:
- Gold cleaning effect: Aqua regia is known to dissolve gold and is also capable of cleaning tarnished or dull gold by removing the surface oxidation. When gold is dipped in this solution, it can appear to shine, as the acid removes any impurities and oxidation layers, giving the gold a bright, fresh appearance.
- Weight loss: The key to solving this puzzle is the drastic reduction in weight. While the surface of the gold may be cleaned and appear shinier, if the solution used is strong enough, it could also dissolve a small amount of the gold itself, causing a reduction in weight. Goldsmiths and fraudsters sometimes use this tactic to reduce the weight of the gold without the victim noticing initially.
- Hasty retreat: The man’s quick departure suggests that he knew he had committed a fraudulent act, and the lady would soon realize the loss of gold.
Thus, the solution used is likely a strong acid that can clean the gold’s surface while causing a reduction in weight, such as aqua regia or nitric acid.
Q 16. Give reasons why copper is used to make hot water tanks and not steel (an alloy of iron).
Ans 16: Copper is used to make hot water tanks instead of steel because it has better thermal conductivity, meaning it heats up and cools down faster. Copper is also more resistant to corrosion and rust compared to steel, especially in the presence of water, making it more durable and efficient for long-term use.
Chapter 3 Metals and Non-metals Class 10 Science Solution Pdf Download
Updated Solution 2024-2025
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