Chemical Reactions and Equations

Daily Life Situations and the Nature of Chemical Change

According to Martin H. Fischer, facts are not science just as the dictionary is not literature. To understand science, we must consider everyday situations and what happens when they occur. Examples include milk being left at room temperature during summers, an iron tawa, pan, or nail being left exposed to a humid atmosphere, grapes getting fermented, food being cooked, food being digested in our body, and the process of respiration. In all these situations, both the nature and the identity of the initial substance have undergone some degree of change. Whenever a chemical change occurs, it can be stated that a chemical reaction has taken place. Determining exactly what a chemical reaction entails and how we recognize its occurrence requires experimental observation.

Experimental Observations: Magnesium, Lead Nitrate, and Zinc

Activity 1.1 involves cleaning a magnesium ribbon about $3-4\,\text{cm}$ long by rubbing it with sandpaper to remove any impurities. The ribbon is held with a pair of tongs and burnt using a spirit lamp or burner. The resulting ash is collected in a watch-glass. It is observed that the magnesium ribbon burns with a dazzling white flame and changes into a white powder. This powder is magnesium oxide ($MgO$), which is formed due to the reaction between magnesium and the oxygen present in the air. During this activity, it is recommended that students wear suitable eyeglasses and keep the ribbon as far from their eyes as possible.

Activity 1.2 involves taking a lead nitrate solution in a test tube and adding a potassium iodide solution to it. Activity 1.3 involves placing a few zinc granules in a conical flask or a test tube and adding dilute hydrochloric acid or sulphuric acid. Observations from this experiment show that bubbles are formed around the zinc granules, indicating the evolution of hydrogen gas ($H_{2}$). Touching the flask reveals a change in temperature. From these three activities, it can be concluded that a chemical reaction is characterized by any of the following: a change in state, a change in colour, the evolution of a gas, or a change in temperature.

Word Equations and Chemical Equations

A chemical reaction can be described in a sentence form, but it is often long and cumbersome. For instance, Activity 1.1 is described as: "when a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide." A shorter representation is a word-equation, such as: $\text{Magnesium} + \text{Oxygen} \rightarrow \text{Magnesium oxide}$ In this equation, magnesium and oxygen are the reactants, performing the chemical change on the left-hand side (LHS). The new substance formed, magnesium oxide, is the product, written on the right-hand side (RHS). Reactants and products are separated by an arrow pointing toward the products, indicating the direction of the reaction. If multiple reactants or products are involved, they are linked by a plus sign ($+$).

A chemical equation can be made even more concise by using chemical formulae instead of words. The word equation for the burning of magnesium is represented symbolically as: $Mg + O_{2} \rightarrow MgO$ When the number of atoms of each element is not the same on both sides of the arrow, the mass is not equal, and the equation is considered unbalanced. Such an equation is called a skeletal chemical equation.

The Law of Conservation of Mass and Balancing Equations

The law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction. Consequently, the total mass of the elements present in the products must be equal to the total mass of the elements present in the reactants. This means the number of atoms of each element must remain the same before and after the reaction. For example, in the reaction between zinc and sulphuric acid: $Zn + H_{2}SO_{4} \rightarrow ZnSO_{4} + H_{2}$ The number of atoms for Zn (1), H (2), S (1), and O (4) are identical on both the LHS and RHS, making it a balanced chemical equation.

The Hit-and-Trial Method of Balancing

To balance a complex skeletal equation like $Fe + H_{2}O \rightarrow Fe_{3}O_{4} + H_{2}$, a step-by-step hit-and-trial method is used. Step I involves drawing boxes around each formula; no changes are made inside these boxes. Step II involves listing the number of atoms: Fe (1 LHS, 3 RHS), H (2 LHS, 2 RHS), and O (1 LHS, 4 RHS). Step III suggests starting with the compound containing the maximum number of atoms. In this case, oxygen in $Fe_{3}O_{4}$ is balanced by putting a coefficient of '4' before $H_{2}O$ on the LHS ($4H_{2}O$). We cannot alter formulae like $H_{2}O_{4}$ or $(H_{2}O)_{4}$. The equation becomes $Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + H_{2}$.

Step IV balances the hydrogen atoms. There are now 8 hydrogen atoms on the LHS, so the coefficient of $H_{2}$ on the RHS is changed to 4, resulting in $Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + 4H_{2}$. Step V balances the iron atoms. There are 3 Fe atoms on the RHS, so a coefficient of 3 is placed before Fe on the LHS. Step VI involves a final count to check the correctness: $3Fe + 4H_{2}O \rightarrow Fe_{3}O_{4} + 4H_{2}$. Both sides have 3 iron, 8 hydrogen, and 4 oxygen atoms. This method is called hit-and-trial because we trial various small whole number coefficients to achieve balance.

Inclusion of Physical States and Reaction Conditions

To make equations more informative, physical states are indicated: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous (a solution in water). The balanced equation for iron and steam is written as: $3Fe(s) + 4H_{2}O(g) \rightarrow Fe_{3}O_{4}(s) + 4H_{2}(g)$ Note that $(g)$ is used with $H_{2}O$ to indicate that water is in the form of steam. Additionally, reaction conditions such as temperature, pressure, or catalysts are written above or below the arrow. For example, the production of methanol occurs at $340\,\text{atm}$ pressure: $CO(g) + 2H_{2}(g) \xrightarrow{340\,\text{atm}} CH_{3}OH(l)$ The photosynthesis reaction is another example, requiring sunlight and chlorophyll: $6CO_{2}(aq) + 12H_{2}O(l) \xrightarrow{\text{Sunlight, Chlorophyll}} C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) + 6H_{2}O(l)$

Combination Reactions and Exothermic Processes

Chemical reactions involve the breaking and making of bonds between atoms. In a combination reaction, two or more reactants combine to form a single product. Activity 1.4 demonstrates this by adding water to calcium oxide (quick lime). The reaction is vigorous and produces slaked lime (calcium hydroxide), releasing a large amount of heat: $CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq) + \text{Heat}$ Reactions where heat is released alongside the formation of products are called exothermic chemical reactions. Other examples include the burning of coal ($C + O_{2} \rightarrow CO_{2}$), the formation of water ($2H_{2} + O_{2} \rightarrow 2H_{2}O$), and the burning of natural gas ($CH_{4} + 2O_{2} \rightarrow CO_{2} + 2H_{2}O$).

Respiration is also an exothermic process. During digestion, carbohydrates in food (like rice and potatoes) are broken down into glucose. Glucose then combines with oxygen in the body's cells to provide energy: $C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) \rightarrow 6CO_{2}(aq) + 6H_{2}O(l) + \text{energy}$ The decomposition of vegetable matter into compost is a further example of an exothermic reaction.

Whitewashing and Calcium Carbonate

A solution of slaked lime ($Ca(OH)_{2}$) is used for whitewashing walls. Over two to three days, the calcium hydroxide reacts slowly with carbon dioxide in the air to form a thin, shiny layer of calcium carbonate ($CaCO_{3}$) and water. It is notably the same chemical formula as marble: $Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)$

Decomposition Reactions

A decomposition reaction occurs when a single reactant breaks down into simpler products. In Activity 1.5, heating green ferrous sulphate crystals ($FeSO_{4} \cdot 7H_{2}O$) causes them to lose water and change colour. They further decompose into solid ferric oxide ($Fe_{2}O_{3}$) and gaseous sulphur dioxide ($SO_{2}$) and sulphur trioxide ($SO_{3}$): $2FeSO_{4}(s) \xrightarrow{\text{Heat}} Fe_{2}O_{3}(s) + SO_{2}(g) + SO_{3}(g)$ Thermal decomposition refers to decomposition carried out by heating. An important industrial example is the decomposition of limestone into quick lime and carbon dioxide, used in cement manufacturing: $CaCO_{3}(s) \xrightarrow{\text{Heat}} CaO(s) + CO_{2}(g)$

Activity 1.6 shows that heating lead nitrate powder in a boiling tube produces brown fumes of nitrogen dioxide ($NO_{2}$): $2Pb(NO_{3})_{2}(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_{2}(g) + O_{2}(g)$ Energy for decomposition can also come from electricity or light. Activity 1.7 details the electrolysis of water, where electric current passed through water in a plastic mug with carbon electrodes produces hydrogen (at the cathode) and oxygen (at the anode). The volume of hydrogen collected is double that of oxygen. In Activity 1.8, white silver chloride ($AgCl$) turns grey in sunlight as it decomposes into silver and chlorine. A similar reaction occurs with silver bromide ($AgBr$), and both are used in black and white photography: $2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_{2}(g)$$2AgBr(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Br_{2}(g)$ Reactions like these, which absorb energy, are called endothermic reactions.

Displacement and Double Displacement Reactions

In a displacement reaction, a more reactive element displaces a less reactive element from its compound. Activity 1.9 shows that an iron nail placed in blue copper sulphate solution becomes brownish, while the blue colour of the solution fades because iron displaces copper: $Fe(s) + CuSO_{4}(aq) \rightarrow FeSO_{4}(aq) + Cu(s)$ Zinc and lead are also more reactive than copper and can displace it from compounds like copper sulphate or copper chloride.

A double displacement reaction involves an exchange of ions between reactants. In Activity 1.10, mixing sodium sulphate and barium chloride solutions forms a white precipitate of barium sulphate ($BaSO_{4}$): $Na_{2}SO_{4}(aq) + BaCl_{2}(aq) \rightarrow BaSO_{4}(s) + 2NaCl(aq)$ Any reaction that produces an insoluble solid (precipitate) is a precipitation reaction. Another example is the reaction between lead(II) nitrate and potassium iodide, which forms a yellow precipitate of lead iodide.

Oxidation and Reduction (Redox)

Oxidation is the gain of oxygen or the loss of hydrogen by a substance. Reduction is the loss of oxygen or the gain of hydrogen. In Activity 1.11, heating copper powder in air produces black copper(II) oxide ($2Cu + O_{2} \rightarrow 2CuO$). If hydrogen gas is passed over the heated $CuO$, it reduces back to brown copper: $CuO + H_{2} \xrightarrow{\text{Heat}} Cu + H_{2}O$ In this reaction, $CuO$ loses oxygen (reduced) and $H_{2}$ gains oxygen (oxidised). Such reactions where one reactant is oxidised and the other is reduced are called redox reactions. Other examples include: $ZnO + C \rightarrow Zn + CO$$MnO_{2} + 4HCl \rightarrow MnCl_{2} + 2H_{2}O + Cl_{2}$

Effects of Oxidation in Everyday Life: Corrosion and Rancidity

Corrosion occurs when a metal is attacked by substances like moisture or acids. Iron undergoes rusting, forming a reddish-brown powder. Silver develops a black coating, and copper develops a green coating. Corrosion causes massive damage to iron structures like bridges, ships, and cars, leading to significant replacement costs annually.

Rancidity occurs when fats and oils are oxidised, changing their smell and taste. To prevent this, antioxidants are added to foods, or food is stored in airtight containers. Chip manufacturers flush bags with nitrogen gas ($N_{2}$) to create an inert environment that prevents oxidation.

Questions & Discussion

Q1: Why should a magnesium ribbon be cleaned before burning in air? Magnesium ribbon is cleaned with sandpaper to remove the layer of magnesium oxide that might have formed on its surface due to reaction with air, ensuring it burns properly.

Q2: Write the balanced equation for the following chemical reactions. (i) Hydrogen + Chlorine $\\rightarrow$ Hydrogen chloride: $H_{2} + Cl_{2} \rightarrow 2HCl$ (ii) Barium chloride + Aluminium sulphate $\\rightarrow$ Barium sulphate + Aluminium chloride: $3BaCl_{2} + Al_{2}(SO_{4})_{3} \rightarrow 3BaSO_{4} + 2AlCl_{3}$ (iii) Sodium + Water $\\rightarrow$ Sodium hydroxide + Hydrogen: $2Na + 2H_{2}O \rightarrow 2NaOH + H_{2}$

Q3: Write balanced chemical equations with state symbols. (i) Barium chloride and sodium sulphate in water: $BaCl_{2}(aq) + Na_{2}SO_{4}(aq) \rightarrow BaSO_{4}(s) + 2NaCl(aq)$ (ii) Sodium hydroxide and hydrochloric acid in water: $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_{2}O(l)$

Q4: A solution of a substance 'X' is used for whitewashing. Name 'X', write its formula, and its reaction with water. 'X' is calcium oxide (quick lime). Formula: $CaO$. Reaction with water: $CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)$.

Q5: Why is the amount of gas collected in one test tube in Activity 1.7 double the other? Water consists of two parts hydrogen and one part oxygen ($2H_{2}O \rightarrow 2H_{2} + O_{2}$). Therefore, the volume of hydrogen gas collected is double the volume of oxygen gas.

Q6: Why does the colour of copper sulphate solution change when an iron nail is dipped in it? Iron is more reactive than copper and displaces copper from the blue copper sulphate solution to form green iron sulphate ($FeSO_{4}$).

Q7: Give an example of a double displacement reaction. Lead nitrate reacts with potassium iodide to form lead iodide and potassium nitrate: $Pb(NO_{3})_{2} + 2KI \rightarrow PbI_{2} + 2KNO_{3}$.

Q8: Identify substances oxidised and reduced in the following: (i) $4Na(s) + O_{2}(g) \rightarrow 2Na_{2}O(s)$: Sodium is oxidised ($Na \rightarrow Na_{2}O$). Oxygen is reduced. (ii) $CuO(s) + H_{2}(g) \rightarrow Cu(s) + H_{2}O(l)$: Copper(II) oxide is reduced to copper; Hydrogen is oxidised to water.