Chemical Reactions and Equations

Decomposition reaction. A single reactant decomposes into two or more simpler products when energy is supplied in the form of heat, light, or electricity.

The brown surface of copper powder develops a black coating of copper(II) oxide ($CuO$) due to oxidation by oxygen in air.

White silver chloride decomposes into grey metallic silver ($Ag$) and yellowish-green chlorine gas ($Cl_2$) in the presence of sunlight (Photolytic Decomposition).

Iron is more reactive than copper in the reactivity series, so it displaces copper from copper sulphate solution. The blue colour of $CuSO_4$ fades to light green (formation of $FeSO_4$) and a reddish-brown copper deposit forms on the iron nail.

• Oxidation: Hydrogen ($H_2$) gains oxygen to form water ($H_2O$) $\implies$ $H_2$ is oxidised (and acts as the Reducing Agent).
• Reduction: Copper(II) oxide ($CuO$) loses oxygen to form metallic copper ($Cu$) $\implies$ $CuO$ is reduced (and acts as the Oxidising Agent).

1. Combination Reaction: Two or more reactants combine to form a single product.
$$CaO (s) + H_2O (l) o Ca(OH)_2 (aq) + Heat$$

2. Decomposition Reaction: A single compound breaks down into two or more simpler substances.
$$CaCO_3 (s) \xr→{\Delta} CaO (s) + CO_2 (g)$$

3. Displacement Reaction: A more reactive element displaces a less reactive element from its compound.
$$Zn (s) + CuSO_4 (aq) o ZnSO_4 (aq) + Cu (s)$$

4. Double Displacement Reaction: Mutual exchange of ions between two compounds.
$$Na_2SO_4 (aq) + BaCl_2 (aq) o BaSO_4 (s)\downarrow + 2NaCl (aq)$$

• Corrosion: The slow, gradual deterioration of metals due to chemical action of moisture, air, and acids (e.g., Rusting of iron: $Fe_2O_3 \cdot xH_2O$).
• Rancidity: Aerial oxidation of fats and oils in food leading to unpleasant smell and taste.
• Prevention methods for iron: (1) Galvanisation (coating with zinc), (2) Painting/oiling/greasing, (3) Alloying (e.g., Stainless steel).

(a) $CH_4 (g) + 2O_2 (g) o CO_2 (g) + 2H_2O (g) + Heat$ $\implies$ Exothermic / Combustion Reaction

(b) $2FeSO_4 (s) \xr→{\Delta} Fe_2O_3 (s) + SO_2 (g) + SO_3 (g)$ $\implies$ Thermal Decomposition Reaction

(c) $BaCl_2 (aq) + Na_2SO_4 (aq) o BaSO_4 (s)\downarrow + 2NaCl (aq)$ $\implies$ Precipitation / Double Displacement Reaction

(d) $2H_2O (l) \xr→{Electricity} 2H_2 (g) + O_2 (g)$ $\implies$ Electrolytic Decomposition Reaction ($2:1$ volume ratio)

(e) $CaO (s) + H_2O (l) o Ca(OH)_2 (aq) + Heat$ $\implies$ Combination / Exothermic Reaction

Yes, as a new substance was formed in the form of a gas.

When baking soda (sodium hydrogen carbonate, $\text{NaHCO}_3$) reacts with vinegar (dilute acetic acid, $\text{CH}_3\text{COOH}$), a brisk effervescence occurs releasing carbon dioxide gas ($\text{CO}_2$), which inflates the balloon. The formation of a chemically new gaseous substance with entirely different properties confirms a chemical change.

Activity 1, as the properties of the substances in the mixture change.

Baking involves endothermic chemical reactions among flour, leavening agents, and sugar that create completely new chemical compounds with irreversible texture, taste, and molecular composition. Slicing pie, crushing a can, and carving wood merely alter physical dimensions without altering chemical identity (physical changes).

reactants - sodium and chlorine; products - sodium chloride

Substances that take part and undergo chemical transformation on the left-hand side of the reaction arrow are reactants ($\text{Na}$ and $\text{Cl}_2$), while the resulting substance formed on the right-hand side is the product ($\text{NaCl}$).

Reactants: $\text{H}$, $\text{Cl}$, $2\text{H}$, and $\text{O}_2$; Products: $\text{HCl}$ and $2\text{H}_2\text{O}$

All chemical entities written to the left of the transformation arrow are reactants; all entities produced to the right of the arrow are products.

$2\text{Al} + 3\text{Cl}_2 \rightarrow 2\text{AlCl}_3$

Aluminium has a valency of $+3$ ($\text{Al}^{3+}$) and chlorine has a valency of $-1$ ($\text{Cl}^-$), so the stable molecular formula of aluminium chloride is $\text{AlCl}_3$. Balancing aluminium metal with diatomic chlorine gas requires $2\text{Al} + 3\text{Cl}_2 \rightarrow 2\text{AlCl}_3$.

$2\text{HCl} + \text{Ba(OH)}_2 \rightarrow \text{BaCl}_2 + 2\text{HOH}$ (or $2\text{H}_2\text{O}$)

Barium hydroxide ($\text{Ba(OH)}_2$, base) neutralizes hydrochloric acid ($\text{HCl}$, acid) to yield barium chloride salt ($\text{BaCl}_2$) and water ($\text{H}_2\text{O}$). Since barium ions ($\text{Ba}^{2+}$) require two chloride ions ($\text{Cl}^-$), $2$ molecules of $\text{HCl}$ must react per formula unit of $\text{Ba(OH)}_2$.

Pb: 1 (LHS), 1 (RHS); Cu: 1 (LHS), 1 (RHS); Cl: 2 (LHS), 2 (RHS)

By atom counting: LHS has $1\text{ Pb}$, $1\text{ Cu}$, and $2\text{ Cl}$. RHS has $1\text{ Pb}$, $1\text{ Cu}$, and $2\text{ Cl}$. The total number of atoms of each element remains strictly conserved across the chemical change.

Na: 2 (LHS), 2 (RHS); Cl: 2 (LHS), 2 (RHS)

LHS contains $2$ atoms of elemental $\text{Na}$ and $1$ diatomic molecule of $\text{Cl}_2$ ($2$ atoms of $\text{Cl}$). RHS contains $2$ formula units of $\text{NaCl}$ ($2$ atoms of $\text{Na}$ and $2$ atoms of $\text{Cl}$).

The reaction is an example of a combination reaction as two compounds react to form a single compound.

Two reactant species (carbon monoxide gas and hydrogen gas) unite chemically under pressure and catalyst to yield one single product (liquid methanol, $\text{CH}_3\text{OH}$). Hence, it is a combination reaction.

Reaction P is an example of a combination reaction while reaction Q is an example of a decomposition reaction

In reaction P, two separate compounds ($\text{CaO} + \text{SO}_2$) combine into a single compound ($\text{CaSO}_3$). In reaction Q, a single reactant compound (zinc carbonate, $\text{ZnCO}_3$) breaks down upon heating into two simpler compounds ($\text{ZnO} + \text{CO}_2$).

A displacement reaction takes place as lead replaces copper from the solution.

In the metal reactivity series, lead ($\text{Pb}$) is more reactive than copper ($\text{Cu}$), so lead displaces copper from $\text{CuSO}_4$, causing the blue colour of the solution to fade and depositing reddish-brown copper. Silver ($\text{Ag}$) is less reactive than copper, giving no reaction.

Silver nitrate and potassium undergo double displacement reaction to form silver chloride and potassium nitrate

A mutual exchange of ions takes place between the two ionic reactants: $\text{Ag}^+$ pairs with $\text{Cl}^-$ to precipitate insoluble white $\text{AgCl}$, while $\text{K}^+$ pairs with $\text{NO}_3^-$ to remain in solution as $\text{KNO}_3$. This is a classic double displacement (precipitation) reaction.

$\text{Zn} \rightarrow \text{Zn}^{2+}$

Oxidation is defined in electronic theory as the loss of electrons (increase in oxidation state). Neutral zinc metal ($\text{Zn}^0$) loses two electrons to form $\text{Zn}^{2+}$, which is an oxidation half-reaction: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$. Conversely, $2\text{H}^+$ gains two electrons to form $\text{H}_2$ (reduction).

Oxidation: $4\text{H}_2$; Reduction: $\text{Fe}_3\text{O}_4$

Hydrogen ($\text{H}_2$) gains oxygen to form $\text{H}_2\text{O}$, so $\text{H}_2$ undergoes oxidation. Magnetic iron oxide ($\text{Fe}_3\text{O}_4$) loses oxygen to form elemental $\text{Fe}$, so $\text{Fe}_3\text{O}_4$ undergoes reduction.

Effect of moisture on metals

Iron articles slowly react with atmospheric oxygen in the presence of moisture (water vapour) to form hydrated ferric oxide (rust, $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$), which is reddish-brown and flaky. This demonstrates the corrosive effect of air and moisture on metals.

The jewellery comes in contact with air, moisture, and acids and corrodes

Silver reacts with trace sulphur compounds (such as hydrogen sulphide, $\text{H}_2\text{S}$) present in air to form a thin black coating of silver sulphide ($\text{Ag}_2\text{S}$):

$$2\text{Ag} (s) + \text{H}_2\text{S} (g) \rightarrow \text{Ag}_2\text{S} (s) [\text{Black}] + \text{H}_2 (g)$$

It prevents rancidity of chips

Potato chips contain fats and oils. When exposed to oxygen in air, fats undergo aerial oxidation into foul-smelling, bad-tasting volatile substances (a process known as rancidity). Flushing sealed packaging with unreactive, inert nitrogen gas ($\text{N}_2$) displaces oxygen and prevents oxidation.

The oils in the bread oxidises and causes rancidity

Food items containing moisture and fats deteriorate through aerobic microbiological action and oxidative rancidity, altering the surface texture and color (development of fungal spores and chemical degradation of organic compounds).

Physical change.

Explanation: Heating iron until it glows red hot only changes its temperature and visual appearance (thermal radiation emission/state of glow) without changing its chemical identity or composition. When cooled back to room temperature, it returns to ordinary metallic iron with no new chemical substance formed.

Three atoms of iron combine with water to form one molecule of iron oxide.

Stoichiometric analysis: 3 atoms (or moles) of iron ($\text{Fe}$) react with 4 molecules of water vapor ($\text{H}_2\text{O}$) to form 1 formula unit/molecule of magnetic iron oxide ($\text{Fe}_3\text{O}_4$) and 4 molecules (8 atoms) of dihydrogen gas ($\text{H}_2$). Water is a molecular substance, not atomic, eliminating options A and B; $\text{Fe}_3\text{O}_4$ is a compound molecule, eliminating option C.

Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) is oxidised.

Explanation: In cellular respiration, glucose loses hydrogen and gains bonds with electronegative oxygen to form carbon dioxide ($\text{CO}_2$). The removal of hydrogen from carbon and addition of oxygen constitutes oxidation. Concurrently, molecular oxygen ($\text{O}_2$) is reduced to water ($\text{H}_2\text{O}$) by the addition of hydrogen.

Products

Explanation: The substances synthesized/formed as a result of a chemical transformation are located on the right side of the reaction arrow and are called products. The starting materials ($ \text{C}_6\text{H}_{12}\text{O}_6 $ and $ \text{O}_2 $) are reactants.

At 1 minute

Reaction Kinetics explanation: The reaction rate is highest at the very start ($t = 0$ to $1$ minute) because the concentrations of the reactants ($[\text{HCl}]$) and available surface area of unreacted magnesium are at their maximum. As the reaction proceeds, reactants are consumed, so the collision frequency decreases and the slope of the gas-volume curve flattens over time.

Reaction rate will increase.

Explanation & Volume at 2 minutes:

• Magnesium powder has a vastly greater specific surface area compared to a single strip of ribbon of identical mass.
• A higher surface area provides far more exposed magnesium atoms for acid particles ($ \text{H}^+ $ ions) to collide with per second, markedly accelerating the reaction rate.
• Therefore, at 2 minutes, a noticeably greater volume of $\text{H}_2$ gas will have already been produced in the measuring cylinder compared to the ribbon experiment.

Correct evaluations:

1. Adding more acid to the flask: No — Increasing the volume without increasing concentration keeps collision frequency per unit area unchanged (assuming acid was already in excess).
2. Heating the acid in the flask: Yes — Higher temperature increases kinetic energy of reactant particles, boosting collision frequency and the fraction of collisions surpassing activation energy.
3. Using a higher concentration of acid: Yes — Greater number of $\text{H}^+$ ions per unit volume leads to higher collision frequency with magnesium surface.

Balanced Chemical Equation:

$$\text{Mg}(s) + 2\text{HCl}(aq) \longrightarrow \text{MgCl}_2(aq) + \text{H}_2(g)$$

Balancing steps: Magnesium has a valency of $+2$ and forms magnesium chloride $\text{MgCl}_2$. Two hydrogen atoms form diatomic gas $\text{H}_2$. Thus, 2 molecules of hydrochloric acid $\text{HCl}$ are required for every 1 atom of magnesium ribbon, conserving all mass.

Rotting of fruits and vegetables

Explanation:

• Rotting/decay: Microorganisms digest and decompose complex organic compounds (polysaccharides, proteins) in dead plant matter into simpler compounds (ammonia, $\text{CO}_2$, $\text{H}_2\text{O}$), which is a microbiological decomposition reaction.
• Melting of glaciers is a physical phase change (solid to liquid).
• Rusting is an oxidation/combination reaction ($4\text{Fe} + 3\text{O}_2 + 2x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$).
• Absorption of $\text{CO}_2$ is dissolution/formation of carbonic acid (combination reaction).

$$\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$$

Combustion reaction: Methane (natural gas fuel) undergoes exothermic oxidation with atmospheric oxygen to yield carbon dioxide gas, water vapor, and liberating substantial thermal energy ($\Delta H < 0$). Mass conservation: 1 C, 4 H, 4 O on both sides.

Exothermic reaction.

Reason: When methane burns in the presence of oxygen, bonds in carbon dioxide and water are formed, releasing large amounts of heat and light energy into the surroundings: $$\text{CH}_4(g) + 2\text{O}_2(g) \longrightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) + \text{Heat}$$

Quick lime is dissolved in water to produce slaked lime.

Explanation: Reaction A (slaking of lime) releases immense thermal energy, causing the water to boil violently ($$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$). Reactions B, C, and D are thermal decomposition reactions that absorb heat energy from the surroundings, making them endothermic.

Complete balanced symbol equation:

$$2\text{Mg}(s) + \text{O}_2(g) \longrightarrow 2\text{MgO}(s)$$

Marking Scheme Breakdown:

• 1 mark: Correct formula for product ($\text{MgO}$).
• 1 mark: Balanced coefficients ($2\text{Mg}$ and $2\text{MgO}$).
• 1 mark: Correct physical state symbols: $(s)$ for $\text{Mg}$, $(g)$ for $\text{O}_2$, and $(s)$ for $\text{MgO}$.

Explanation:

It is classified as a combination reaction because two distinct starting reactants (magnesium metal, an element, and oxygen gas, an element) combine chemically to yield only a single chemical product, magnesium oxide ($\text{MgO}$).

Explanation:

The reaction is an oxidation reaction because magnesium gains oxygen to form magnesium oxide ($\text{MgO}$). In terms of electronic concept, each neutral magnesium atom loses two electrons ($\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-$), which defines oxidation.

Explanation:

The reaction releases energy into the surroundings in the form of intense heat and a dazzling white light; any reaction that releases thermal energy to the surroundings is termed an exothermic reaction.

Key reasons:

1. Magnesium is a reactive metal that slowly reacts with atmospheric oxygen and moisture at room temperature to form a tough, unreactive protective layer of basic magnesium oxide/carbonate ($\text{MgO} / \text{MgCO}_3$) on its surface.
2. Cleaning the ribbon with sandpaper strips off this dull inert coating so that pure underlying metallic magnesium is directly exposed to flame and oxygen, allowing it to ignite smoothly.

Three essential lab precautions:

1. Hold with crucible tongs: Never hold the burning magnesium ribbon with bare hands or forceps that conduct heat; hold it firmly using metal crucible tongs over a watch glass.
2. Wear safety goggles / UV eye protection: Do not look directly at the dazzling, brilliant white flame, as the intense ultraviolet emission can injure the retina.
3. Keep at a safe distance from face and body: Hold the burning ribbon away from clothing and face, placing a watch glass below to safely collect hot falling white ash.

Suggested pH & Explanation:

pH: Approximately 9 to 10 (greater than 7).

Explanation: Magnesium oxide is a metallic oxide, and basic in nature. When mixed with water, it sparingly dissolves to form magnesium hydroxide, a weak base/alkali: $$\text{MgO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Mg(OH)}_2(aq)$$ This releases hydroxide ions ($\text{OH}^-$) into the solution, turning red litmus paper blue and producing an alkaline $\text{pH} > 7$.

a combination reaction

Definition: Reactions in which two or more simple reactants (elements or compounds) combine chemically to produce a single chemical compound ($A + B \rightarrow AB$) are known as combination (or synthesis) reactions.

Electrolytic decomposition of water:

• Definition: Electrolytic decomposition is a chemical reaction where a chemical compound is split into simpler constituent substances by passing an electric current through its molten or aqueous state.
• Chemical Equation: $$2\text{H}_2\text{O}(l) \xrightarrow{\text{Electric current}} 2\text{H}_2(g) \uparrow + \text{O}_2(g) \uparrow$$
• Cathode reaction: Dihydrogen gas ($\text{H}_2$) is liberated at the negative electrode (cathode) by reduction ($2\text{H}^+ + 2e^- \rightarrow \text{H}_2$).
• Anode reaction: Dioxygen gas ($\text{O}_2$) is liberated at the positive electrode (anode) by oxidation ($2\text{O}^{2-} \rightarrow \text{O}_2 + 4e^-$).
• Volume Ratio: Since two parts of hydrogen combine with one part of oxygen by volume in water, the volume of $\text{H}_2$ gas collected at the cathode is exactly twice ($2:1$) the volume of $\text{O}_2$ gas collected at the anode.

Reason: Pure distilled water consists almost entirely of neutral covalent molecules and undergoes negligible self-ionization ($[\text{H}^+] = [\text{OH}^-] = 10^{-7}\text{ M}$). Because it contains virtually no free mobile charge carriers (ions) to conduct electric current, it is an insulator. A few drops of dilute sulfuric acid or an ionic salt must be added to acidulate water and facilitate electrolysis.

Displacement reaction (also a Redox reaction)

Explanation: In this reaction, the more electronegative oxygen atoms displace the hydrogen atoms from ammonia ($\text{NH}_3$), binding with nitrogen to form nitric oxide ($\text{NO}$) and with hydrogen to form water ($\text{H}_2\text{O}$). Nitrogen is oxidized (oxidation state $-3 \rightarrow +2$) and oxygen is reduced ($0 \rightarrow -2$). Hence it is classified as a displacement/redox reaction.

Balanced Symbol Equation:

$$\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$

Marking criteria: 1 mark for correct chemical formulas of calcium oxide ($\text{CaO}$), water ($\text{H}_2\text{O}$), and calcium hydroxide / slaked lime ($\text{Ca(OH)}_2$); 1 mark for stoichiometric balance and state symbols.

Reaction Type: Combination reaction and highly exothermic reaction.

Explanation: Two distinct compounds (calcium oxide and water) react chemically to yield a single compound (calcium hydroxide), satisfying the definition of a combination reaction. In addition, it liberates an enormous quantity of heat energy, raising temperature to boiling levels, making it exothermic.

Two lab precautions:

1. Avoid touching the reaction vessel directly: The beaker becomes intensely hot due to vigorous exothermic heat liberation and can cause severe thermal burns.
2. Add water slowly / wear eye goggles: The vigorous bubbling, steam generation, and spitting of caustic alkaline slaked lime suspension can cause skin and eye irritation.

Identifications:

• Reacting gas: Carbon dioxide ($\text{CO}_2$).
• Other product formed: Water ($\text{H}_2\text{O}$).

Chemical Equation: $$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s) \downarrow + \text{H}_2\text{O}(l)$$ This reaction is used for whitewashing walls, forming a thin, lustrous layer of calcium carbonate after 2 to 3 days, and is also the standard confirmatory test for carbon dioxide (turning lime water milky).

1 and 4 (Iron is oxidised; Water is oxidising agent)

Redox Analysis:

• Iron ($\text{Fe}$): Gains oxygen to form $\text{Fe}_3\text{O}_4$. Hence, Iron is oxidised (Statement 1 is correct).
• Water ($\text{H}_2\text{O}$): Loses oxygen to become hydrogen gas ($\text{H}_2$). Hence, water is reduced (Statement 2 is incorrect).
• Because water provides oxygen and causes the oxidation of iron while being itself reduced, water acts as the oxidising agent (Statement 4 is correct; Statement 3 is incorrect).
• Therefore, statements 1 and 4 are correct.