CBSE Class 10 Science: Metals and Non-metals

Introduction to Metals and Non-metals

Elements are the basic units of matter, and for convenience, chemists broadly classify them into two main categories: metals and non-metals. This classification is primarily based on their physical and chemical properties. Metals, like iron, copper, and gold, are typically solid, lustrous, and good conductors of heat and electricity. Non-metals, such as oxygen, carbon, and sulfur, usually lack these properties and can exist as solids, liquids, or gases. However, this distinction isn't always sharp, and some elements exhibit properties of both, known as metalloids (e.g., silicon, germanium). Understanding these classifications is crucial for predicting how elements will behave in chemical reactions and for understanding their vast applications in technology, industry, and biological systems. For instance, the use of metals in construction and wiring, or non-metals in medicines and fertilizers, are direct consequences of their unique properties.

Physical Properties: Metals vs. Non-metals

Lustre

Metals: Generally shiny (metallic lustre). Non-metals: Usually dull, non-lustrous (exceptions: iodine, graphite).

Hardness

Metals: Generally hard (exceptions: sodium, potassium, lead, which are soft and can be cut with a knife). Non-metals: Mostly soft (exception: diamond, an allotrope of carbon, is the hardest natural substance).

State at Room Temperature

Metals: Solids (exception: mercury is a liquid). Non-metals: Solids, liquids (bromine), or gases.

Malleability

Metals: Malleable (can be beaten into thin sheets) - e.g., gold, silver, aluminium. Non-metals: Non-malleable, brittle (they break when hammered).

Ductility

Metals: Ductile (can be drawn into thin wires) - e.g., gold, copper. Non-metals: Non-ductile.

Conductivity (Heat & Electricity)

Metals: Good conductors of heat and electricity (exceptions: lead, mercury are relatively poor conductors). Non-metals: Poor conductors of heat and electricity (exception: graphite, an allotrope of carbon, is a good conductor of electricity).

Sonority

Metals: Sonorous (produce a ringing sound when struck). Non-metals: Non-sonorous.

Melting and Boiling Points

Metals: Generally high (exceptions: sodium, potassium, mercury). Non-metals: Generally low (exception: diamond, graphite have very high melting points).

Density

Metals: Generally high. Non-metals: Generally low.

Chemical Properties of Metals

1. Reaction with Oxygen/Air — Most metals react with oxygen to form metal oxides. These oxides are generally basic in nature (they turn red litmus blue) and react with water to form bases. Some metal oxides, like aluminium oxide and zinc oxide, are amphoteric, meaning they show both acidic and basic properties. Example: $4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$ (Sodium oxide) Example: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$ (Magnesium oxide) Amphoteric Oxide Example: $Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)$ (Basic nature) Amphoteric Oxide Example: $Al_2O_3(s) + 2NaOH(aq) \rightarrow 2NaAlO_2(aq) + H_2O(l)$ (Acidic nature, forms sodium aluminate)
2. Reaction with Water — Different metals react with water under different conditions (cold water, hot water, or steam) to form metal oxides or hydroxides, and hydrogen gas. Highly reactive metals (Na, K, Ca): React vigorously with cold water to form hydroxides and hydrogen gas. Example: $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g) + Heat$ Less reactive metals (Mg, Al, Zn, Fe): Magnesium reacts with hot water; Aluminium, Zinc, and Iron react with steam to form metal oxides and hydrogen gas. Example: $Mg(s) + 2H_2O(hot) \rightarrow Mg(OH)_2(aq) + H_2(g)$ Example: $2Al(s) + 3H_2O(steam) \rightarrow Al_2O_3(s) + 3H_2(g)$ Least reactive metals (Pb, Cu, Ag, Au): Do not react with water at all.
3. Reaction with Acids — Metals generally react with dilute acids to form metal salts and hydrogen gas. However, the reactivity varies. Example (Dilute HCl): $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$ Example (Dilute $H_2SO_4$): $Fe(s) + H_2SO_4(aq) \rightarrow FeSO_4(aq) + H_2(g)$ * Note: Metals usually do not react with nitric acid to produce hydrogen gas, as nitric acid is a strong oxidising agent and oxidises the $H_2$ produced to water. Magnesium ($Mg$) and Manganese ($Mn$) are exceptions, reacting with very dilute nitric acid to produce $H_2$ gas.
4. Reaction with Salt Solutions of Other Metals — A more reactive metal can displace a less reactive metal from its salt solution. This is known as a displacement reaction, which we'll discuss further. * Example: $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$ (Iron displaces copper, as iron is more reactive than copper).

Chemical Properties of Non-metals

1. Reaction with Oxygen/Air — Non-metals react with oxygen to form non-metallic oxides, which are generally acidic in nature (they turn blue litmus red) or neutral. They do not typically form basic oxides. Acidic Oxide Example: $C(s) + O_2(g) \rightarrow CO_2(g)$ (Carbon dioxide, acidic) Acidic Oxide Example: $S(s) + O_2(g) \rightarrow SO_2(g)$ (Sulfur dioxide, acidic) * Neutral Oxide Example: $N_2(g) + O_2(g) \rightarrow 2NO(g)$ (Nitric oxide, neutral)
2. Reaction with Water — Non-metals generally do not react with water or steam. For example, carbon, sulfur, and phosphorus do not react with water.
3. Reaction with Acids — Non-metals generally do not react with dilute acids. They might react with concentrated oxidising acids (like concentrated $H_2SO_4$ or $HNO_3$) to form their respective oxides or oxyacids.
4. Reaction with Salt Solutions of Other Non-metals — A more reactive non-metal can displace a less reactive non-metal from its salt solution (or halide solution). For example, chlorine can displace bromine from potassium bromide solution. * Example: $Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)$ (Chlorine displaces bromine, as chlorine is more reactive than bromine).

Reactivity Series and Types of Bonds

The Reactivity Series (or Activity Series) is a list of metals arranged in the decreasing order of their reactivity. A more reactive metal can displace a less reactive metal from its salt solution. This series helps predict the outcome of displacement reactions.

Common Reactivity Series (Decreasing Reactivity):
Potassium ($K$) > Sodium ($Na$) > Calcium ($Ca$) > Magnesium ($Mg$) > Aluminium ($Al$) > Zinc ($Zn$) > Iron ($Fe$) > Lead ($Pb$) > Hydrogen ($H$) > Copper ($Cu$) > Mercury ($Hg$) > Silver ($Ag$) > Gold ($Au$)

Elements above hydrogen in the series can displace hydrogen from dilute acids, while those below cannot.

How do elements form bonds?

1. Ionic Bonding: This type of bonding occurs between a metal and a non-metal. Metals tend to lose electrons to achieve a stable noble gas configuration, forming positively charged ions (cations). Non-metals tend to gain these electrons to achieve a stable configuration, forming negatively charged ions (anions). The strong electrostatic force of attraction between these oppositely charged ions forms an ionic bond. For example, in Sodium Chloride ($NaCl$), sodium (metal) loses one electron to become $Na^+$ and chlorine (non-metal) gains that electron to become $Cl^-$.

1. Covalent Bonding: This type of bonding typically occurs between two non-metals. Neither atom completely transfers electrons; instead, they share electrons to achieve a stable noble gas configuration. This sharing of electrons forms a covalent bond. For example, in a water molecule ($H_2O$), hydrogen and oxygen atoms share electrons. Carbon compounds, which are largely covalent, will be studied in detail in a later chapter.

Exam Tip: Avoiding Common Mistakes

1. Exceptions are Key: Remember exceptions to physical properties (e.g., Mercury-liquid metal, Sodium/Potassium-soft metals, Graphite-conductor non-metal, Diamond-hardest non-metal). These are frequent question types.
2. Balancing Equations: Always ensure chemical equations are balanced. Incorrectly balanced equations can lead to loss of marks, especially in chemical properties sections.
3. Reactivity Series: Memorise the reactivity series accurately. Many questions involve predicting displacement reactions based on it. Understand that a metal higher in the series will displace one lower down.
4. Nature of Oxides: Differentiate between basic, acidic, and amphoteric oxides. Metal oxides are generally basic, non-metal oxides are generally acidic. Aluminium and Zinc oxides are common examples of amphoteric oxides.
5. Reaction with Acids (Nitric Acid): Be careful with nitric acid ($HNO_3$). Though an acid, it's a strong oxidising agent. While most metals react with dilute acids to produce hydrogen, with nitric acid, hydrogen is usually oxidised to water. Remember the exceptions: Magnesium ($Mg$) and Manganese ($Mn$) react with very dilute nitric acid to produce hydrogen gas.

Practice Questions with Solutions

• Q: Give reasons for the following: a) Sodium is stored in kerosene oil. b) Aluminium oxide is considered an amphoteric oxide. c) Generally, non-metals do not react with dilute acids. A: Step 1: Address part a). Sodium is a highly reactive metal. It reacts vigorously with oxygen and water present in the air, releasing a large amount of heat and sometimes catching fire. To prevent this reaction and accidental fire, it is stored immersed in kerosene oil, which cuts off its contact with air and moisture. Step 2: Address part b). Aluminium oxide ($Al_2O_3$) reacts with both acids and bases to produce salt and water, exhibiting both basic and acidic properties. Therefore, it is considered an amphoteric oxide. Example with acid: $Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)$ Example with base: $Al_2O_3(s) + 2NaOH(aq) \rightarrow 2NaAlO_2(aq) + H_2O(l)$ Step 3: Address part c). Non-metals are electron acceptors. They tend to gain electrons to complete their outermost shell. When reacting with dilute acids, hydrogen ions ($H^+$) need to gain electrons to form hydrogen gas ($H_2$). Non-metals cannot provide electrons to the $H^+$ ions, and thus generally do not displace hydrogen from acids. Also, non-metals do not typically act as reducing agents required to reduce $H^+$ to $H_2$. Final answer: Sodium is stored in kerosene due to its high reactivity. Aluminium oxide is amphoteric as it reacts with both acids and bases. Non-metals do not react with dilute acids because they are electron acceptors and cannot donate electrons to displace hydrogen.
• Q: Write balanced chemical equations for the following reactions: a) Iron reacts with steam. b) Magnesium reacts with dilute hydrochloric acid. c) Chlorine gas reacts with potassium bromide solution. A: Step 1: Write the equation for Iron reacting with steam. Iron reacts with steam to form iron(II,III) oxide (magnetite) and hydrogen gas. $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$ Step 2: Write the equation for Magnesium reacting with dilute hydrochloric acid. Magnesium reacts with dilute hydrochloric acid to form magnesium chloride and hydrogen gas. $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$ Step 3: Write the equation for Chlorine gas reacting with potassium bromide solution. Chlorine, being more reactive than bromine, displaces bromine from potassium bromide solution to form potassium chloride and bromine. $Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)$ Final answer: The balanced equations are: a) $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$ b) $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$ c) $Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)$.
• Q: A student took a clean iron nail and placed it in a test tube containing copper sulphate solution. After about 20 minutes, he observed the solution and the iron nail. What changes would he observe? Write the chemical equation for the reaction. A: Step 1: Identify the reactants and their relative reactivities. The reactants are iron (Fe) and copper sulphate solution ($CuSO_4$). According to the reactivity series, iron is more reactive than copper. Step 2: Predict the type of reaction and observations. Since iron is more reactive than copper, iron will displace copper from its salt solution. The blue colour of copper sulphate solution will fade and eventually turn light green due to the formation of iron(II) sulphate. A reddish-brown deposit of copper metal will be observed on the iron nail. Step 3: Write the balanced chemical equation. $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$ Final answer: The student would observe that the blue colour of the copper sulphate solution changes to light green, and a reddish-brown deposit would form on the iron nail. The chemical equation is $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$.
• Q: Describe the formation of an ionic compound with the help of an example. Mention any two general properties of ionic compounds. A: Step 1: Describe ionic compound formation with an example. Ionic compounds are formed when electrons are transferred from a metal atom to a non-metal atom, resulting in the formation of positively charged ions (cations) and negatively charged ions (anions). These oppositely charged ions are held together by strong electrostatic forces of attraction, forming an ionic bond. Example: Formation of Magnesium Chloride ($MgCl_2$). Magnesium (Mg) is a metal with atomic number 12 (electronic configuration: 2, 8, 2). It tends to lose its two valence electrons to achieve a stable noble gas configuration, forming an $Mg^{2+}$ ion. Chlorine (Cl) is a non-metal with atomic number 17 (electronic configuration: 2, 8, 7). It needs one electron to achieve a stable configuration, forming a $Cl^-$ ion. Since magnesium loses two electrons, two chlorine atoms are required to accept one electron each. $Mg \rightarrow Mg^{2+} + 2e^-$ $2Cl + 2e^- \rightarrow 2Cl^-$ The $Mg^{2+}$ and $2Cl^-$ ions then attract each other to form $MgCl_2$. Step 2: Mention two general properties of ionic compounds. 1. Physical Nature: Ionic compounds are generally solid and hard due to the strong electrostatic forces of attraction between the ions. They are often brittle. 2. Melting and Boiling Points: They have high melting and boiling points because a considerable amount of energy is required to break the strong inter-ionic attraction. 3. Solubility: Most ionic compounds are soluble in water but insoluble in organic solvents like kerosene or petrol. 4. Conductivity: They conduct electricity in molten state and in aqueous solutions, but not in the solid state. In molten or aqueous state, the ions are free to move and conduct current. Final answer: Ionic compounds form via electron transfer from metal to non-metal, creating charged ions that attract (e.g., $MgCl_2$). Two properties are high melting/boiling points and conductivity in molten/aqueous states.
• Q: What is thermite reaction? Give its balanced chemical equation and mention one of its applications. A: Step 1: Define thermite reaction. Thermite reaction is a highly exothermic redox reaction between a metal oxide and a more reactive metal (usually aluminium powder) which acts as a reducing agent. It produces a large amount of heat, melting the metal product. Step 2: Write the balanced chemical equation. One common example involves iron(III) oxide and aluminium: $Fe_2O_3(s) + 2Al(s) \xrightarrow{Heat} 2Fe(l) + Al_2O_3(s) + Heat$ In this reaction, aluminium is more reactive than iron, so it displaces iron from its oxide. The iron produced is in a molten state due to the immense heat generated. Step 3: Mention one application. Thermite reactions are used for welding railway tracks or cracked machine parts, as the molten iron produced can flow into the cracks and fuse the pieces together. Final answer: Thermite reaction is a highly exothermic redox reaction, usually between a metal oxide and aluminium powder. The balanced equation for iron(III) oxide is $Fe_2O_3(s) + 2Al(s) \xrightarrow{Heat} 2Fe(l) + Al_2O_3(s) + Heat$. It is used for welding railway tracks or cracked machine parts.

Frequently Asked Questions