Chemical Reactions and Equations

A chemical reaction involves the transformation of reactants into products through the breaking and forming of bonds. It can be identified by changes such as evolution of gas, change in temperature, formation of precipitate, or change in color. For example, burning magnesium ribbon in air produces magnesium oxide, indicated by a dazzling white flame and a white powder.

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. This means the total mass of reactants equals the total mass of products. For instance, in the reaction 2H2 + O2 → 2H2O, the mass of hydrogen and oxygen reactants equals the mass of water produced.

Balancing a chemical equation involves ensuring the same number of each type of atom on both sides of the equation. Start by listing the number of atoms of each element on both sides, then adjust coefficients to balance them. For example, to balance Fe + H2O → Fe3O4 + H2, you would adjust coefficients to get 3Fe + 4H2O → Fe3O4 + 4H2.

A combination reaction is where two or more substances combine to form a single product. An example is the formation of water from hydrogen and oxygen: 2H2 + O2 → 2H2O. These reactions often release energy, making them exothermic.

A decomposition reaction occurs when a single compound breaks down into two or more simpler substances. For example, heating calcium carbonate decomposes it into calcium oxide and carbon dioxide: CaCO3 → CaO + CO2. These reactions usually require energy input, making them endothermic.

Exothermic reactions release energy, often as heat, like the combustion of methane: CH4 + 2O2 → CO2 + 2H2O + energy. Endothermic reactions absorb energy, such as the decomposition of water into hydrogen and oxygen using electricity: 2H2O → 2H2 + O2.

A displacement reaction involves an element displacing another in a compound. For example, iron displaces copper in copper sulphate solution: Fe + CuSO4 → FeSO4 + Cu. The more reactive element (iron) displaces the less reactive one (copper).

Double displacement reactions involve the exchange of ions between two compounds, often forming a precipitate. An example is the reaction between sodium sulphate and barium chloride: Na2SO4 + BaCl2 → BaSO4 + 2NaCl, where barium sulphate is the precipitate.

Oxidation is the gain of oxygen or loss of hydrogen, while reduction is the loss of oxygen or gain of hydrogen. In the reaction CuO + H2 → Cu + H2O, CuO is reduced to Cu, and H2 is oxidized to H2O. These processes often occur together in redox reactions.

Respiration is exothermic because it releases energy by breaking down glucose in the presence of oxygen: C6H12O6 + 6O2 → 6CO2 + 6H2O + energy. This energy is used by cells for various functions, making it a vital biological process.

Corrosion is the deterioration of metals due to reactions with environmental substances like oxygen and moisture. For example, iron rusts when exposed to moist air, forming iron oxide. This process weakens the metal and causes significant economic losses.

Rancidity can be prevented by storing food in airtight containers, using antioxidants, or flushing with inert gases like nitrogen. These methods slow down oxidation, which causes fats and oils to develop unpleasant smells and tastes.

A precipitation reaction forms an insoluble solid (precipitate) when two solutions are mixed. For example, mixing silver nitrate and sodium chloride solutions produces silver chloride precipitate: AgNO3 + NaCl → AgCl + NaNO3.

Magnesium ribbon is cleaned to remove the oxide layer that forms on its surface, ensuring a pure reaction with oxygen. This allows the magnesium to burn more efficiently, producing a bright white flame and forming magnesium oxide.

Zinc reacts with dilute sulphuric acid to produce zinc sulphate and hydrogen gas: Zn + H2SO4 → ZnSO4 + H2. The reaction is exothermic, releasing heat, and the hydrogen gas can be tested by bringing a burning candle near it, producing a pop sound.

When lead nitrate reacts with potassium iodide, a double displacement reaction occurs, forming lead iodide precipitate and potassium nitrate: Pb(NO3)2 + 2KI → PbI2 + 2KNO3. The yellow precipitate of lead iodide indicates the reaction.

Chemical equations provide a concise representation of reactions, showing reactants, products, and their stoichiometry. They help predict reaction outcomes, balance mass, and understand energy changes, making them fundamental in chemical studies and industries.

Electrolysis of water decomposes it into hydrogen and oxygen gases using electricity: 2H2O → 2H2 + O2. This reaction shows decomposition as a single compound (water) breaks down into two simpler substances (hydrogen and oxygen).

Catalysts speed up reactions without being consumed. For example, in the decomposition of hydrogen peroxide, manganese dioxide acts as a catalyst: 2H2O2 → 2H2O + O2. Catalysts lower activation energy, making reactions faster and more efficient.

The reaction between calcium oxide and water is exothermic because it releases heat: CaO + H2O → Ca(OH)2 + heat. This energy release is due to the formation of new bonds in calcium hydroxide, making the reaction highly energetic.

When silver chloride is exposed to sunlight, it decomposes into silver and chlorine gas: 2AgCl → 2Ag + Cl2. This photodecomposition reaction is used in black and white photography to capture images.

Iron displaces copper in copper sulphate solution because it is more reactive: Fe + CuSO4 → FeSO4 + Cu. The blue color of copper sulphate fades as iron sulphate forms, and copper deposits on the iron, showing a classic displacement reaction.

Thermal decomposition uses heat to break down compounds, like calcium carbonate into calcium oxide and CO2. Photochemical decomposition uses light, such as silver chloride breaking into silver and chlorine under sunlight. Both are decomposition but differ in energy sources.

The reaction BaCl2 + Na2SO4 → BaSO4 + 2NaCl is double displacement because the ions of the reactants exchange partners. Barium pairs with sulphate to form a precipitate, and sodium pairs with chloride, remaining in solution.

Burning hydrogen sulphide in air produces water and sulphur dioxide: 2H2S + 3O2 → 2H2O + 2SO2. This combustion reaction is exothermic, releasing energy, and is notable for the pungent smell of sulphur dioxide.