Chemical Reaction - Class 10 AP Board General Science

Chemical Reactions and Equations

Physical and Chemical Changes

Physical changes involve alterations in state, shape, size, or color without changing the chemical composition of the substance. Chemical changes result in the formation of new substances with different chemical properties and are usually permanent and non-reversible.

• Physical Change: No new substance formed; reversible; examples include melting of ice, boiling of water, folding paper.
• Chemical Change: New substance formed; irreversible; examples include burning paper, fermentation of grapes to wine, milk turning to curd, rusting of iron.

Chemical Reactions: Definition and Features

A chemical reaction is a process…

Chemical Reactions and Equations

Physical and Chemical Changes

Physical changes involve alterations in state, shape, size, or color without changing the chemical composition of the substance. Chemical changes result in the formation of new substances with different chemical properties and are usually permanent and non-reversible.

• Physical Change: No new substance formed; reversible; examples include melting of ice, boiling of water, folding paper.
• Chemical Change: New substance formed; irreversible; examples include burning paper, fermentation of grapes to wine, milk turning to curd, rusting of iron.

Chemical Reactions: Definition and Features

A chemical reaction is a process in which new chemical substances are formed from reactants. The process involves changes at the molecular level, resulting in products with properties different from the reactants.

• Key features of chemical reactions:
  • Formation of new substances.
  • Permanent and non-reversible changes.
  • Observable changes such as color, temperature, state, evolution of gas, or formation of precipitate.

Writing and Balancing Chemical Equations

Chemical equations represent chemical reactions using symbols and formulas. Reactants are written on the left, products on the right, separated by an arrow. Balancing ensures the number of atoms of each element is equal on both sides, adhering to the law of conservation of mass.

• Symbols for physical states:
  • $(s)$: solid
  • $(l)$: liquid
  • $(g)$: gas
  • $(aq)$: aqueous solution
• Balancing method:
  • Do not change subscripts in formulas.
  • Adjust coefficients in front of compounds/elements.
  • Balance metals first, then non-metals, then oxygen, then hydrogen.
• Example: $2\ \text{Mg} + \text{O}_2 \rightarrow 2\ \text{MgO}$

Types of Chemical Reactions

Chemical reactions are classified based on the nature of the reactants and products. The main types are:

• Combination Reaction: Two or more reactants combine to form a single product.
  • Example: $\text{Mg} + \text{O}_2 \rightarrow \text{MgO}$
• Decomposition Reaction: A single compound breaks down into two or more simpler substances.
  • Example: $\text{CaCO}_3 \xrightarrow{\text{heat}} \text{CaO} + \text{CO}_2$
• Displacement Reaction: A more reactive element displaces a less reactive element from its compound.
  • Example: $\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$
• Double Displacement Reaction: Exchange of ions between two compounds to form new compounds.
  • Example: $\text{BaCl}_2 + \text{Na}_2\text{SO}_4 \rightarrow \text{BaSO}_4 + 2\ \text{NaCl}$
• Redox Reaction: Both oxidation and reduction occur simultaneously.

Identification of Reaction Types and Key Examples

Each reaction type has characteristic observations and examples. Recognizing these helps in predicting products and balancing equations.

• Combination: Single product, often exothermic, e.g., burning magnesium ribbon.
• Decomposition: Requires energy input (heat, light, electricity), e.g., heating calcium carbonate.
• Displacement: Occurs in solutions, color changes, e.g., iron nail in copper sulfate turns solution green.
• Double Displacement: Formation of precipitate or gas, e.g., white precipitate of barium sulfate.
• Redox: Involves transfer of oxygen, hydrogen, or electrons.

Valency, Ions, and Compound Formation

Understanding valency and the charges on ions is essential for writing correct chemical formulas. Cations (positive ions) and anions (negative ions) combine in ratios that balance overall charge.

• Common cations: Na$^+$, K$^+$, Ca$^{2+}$, Al$^{3+}$, Fe$^{2+}$, Fe$^{3+}$, Cu$^+$, Cu$^{2+}$
• Common anions: Cl$^-$, Br$^-$, I$^-$, O$^{2-}$, S$^{2-}$, NO$_3^-$, SO$_4^{2-}$, CO$_3^{2-}$
• Naming: Single negative ions end with "-ide" (e.g., chloride, oxide); polyatomic ions have specific names (e.g., sulfate, nitrate).
• Variable valency: Some elements (e.g., copper, iron) show more than one valency, leading to compounds like FeCl$_2$ (ferrous chloride) and FeCl$_3$ (ferric chloride).

Law of Conservation of Mass and Balancing Rationale

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. Therefore, the total mass of reactants equals the total mass of products, and equations must be balanced to reflect this.

• Reason for balancing: Ensures the number of atoms of each element is the same on both sides.
• Example: If 10 kg of reactants are used, 10 kg of products are formed (in a closed system).

Observations During Chemical Reactions

Several observable changes indicate a chemical reaction has occurred. These include:

• Change in state: Solid to liquid, liquid to gas, etc.
• Change in temperature: Heat may be absorbed (endothermic) or released (exothermic).
• Change in color: E.g., green to white, blue to green.
• Evolution of gas: Bubbles or effervescence.
• Formation of precipitate: Insoluble solid settles out.

Mnemonic: FACTS (Formation of precipitate, Appearance of gas, Change in color, Temperature change, State change).

Special Reaction Types: Exothermic, Endothermic, Redox, and Corrosion

• Exothermic reactions: Release heat (e.g., combination reactions, respiration, burning).
• Endothermic reactions: Absorb heat (e.g., decomposition reactions like heating CaCO$_3$).
• Redox reactions: Involve both oxidation (addition of oxygen/removal of hydrogen) and reduction (removal of oxygen/addition of hydrogen).
• Corrosion: Deterioration of metals due to reaction with air and moisture (e.g., rusting of iron, tarnishing of silver, green coating on copper).
• Prevention: Use of anti-oxidants, air-tight containers, or inert gases like nitrogen for food preservation.

Practical Examples and Exam-Focused Reactions

• Burning magnesium ribbon: $2\ \text{Mg} + \text{O}_2 \rightarrow 2\ \text{MgO}$ (white powder, bright light, exothermic)
• Formation of slaked lime: $\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2$ (used in cement, exothermic)
• Testing for hydrogen gas: Burns with a 'pop' sound.
• Testing for oxygen gas: Rekindles a glowing splint.
• Electrolysis of water: $2\ \text{H}_2\text{O} \xrightarrow{\text{electricity}} 2\ \text{H}_2 + \text{O}_2$ (hydrogen collected at cathode, oxygen at anode; hydrogen volume is double that of oxygen).
• Photolytic decomposition: Silver chloride turns grey in sunlight due to decomposition into silver and chlorine gas.

Key Exam Pitfalls and Tips

• Do not alter subscripts in chemical formulas when balancing equations; only adjust coefficients.
• Always balance metals first, then non-metals, then oxygen, then hydrogen.
• Recognize color changes and precipitate formation as evidence of specific reaction types.
• Remember that not all reactions are reversible; chemical changes are generally permanent.
• For variable valency elements, use correct nomenclature (e.g., ferrous vs. ferric).
• In displacement reactions, check the reactivity series to predict if a reaction will occur.
• In double displacement, look for formation of insoluble salts (precipitates) as confirmation.

Table: Common Color Changes in Reactions

• Use this table to quickly recall color changes for common exam reactions.

End of Notes