Chemical changes Revision Notes

Acids and Metals

When an acid reacts with a metal more reactive than hydrogen, a salt and hydrogen gas are produced. The general equation is: Acid + Metal → Salt + Hydrogen. For example, sulfuric acid + zinc → zinc sulfate + hydrogen. Metals less reactive than hydrogen (like copper, mercury, silver, gold, platinum) generally do not react with dilute acids to produce hydrogen.

Acids and Metal Oxides/Hydroxides

Metal oxides and metal hydroxides are bases. When an acid reacts with a base, a neutralisation reaction occurs, producing a salt and water. The general equation is: Acid + Base → Salt + Water. For example, nitric acid + sodium hydroxide → sodium nitrate + water. This type of reaction is exothermic, releasing heat.

Acids and Metal Carbonates

When an acid reacts with a metal carbonate, a salt, water, and carbon dioxide gas are produced. The general equation is: Acid + Metal Carbonate → Salt + Water + Carbon Dioxide. For example, hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide. The carbon dioxide gas causes effervescence.

Acids as Sources of Hydrogen Ions

Acids are substances that, when dissolved in a solution (typically water), produce hydrogen ions (H⁺). These H⁺ ions are responsible for the acidic nature of the solution. For instance, hydrochloric acid (HCl) dissociates in water to form H⁺ and Cl⁻ ions. The presence of these hydrogen ions is what gives acids their characteristic properties, such as turning blue litmus paper red and reacting with carbonates to produce carbon dioxide.

Alkalis as Sources of Hydroxide Ions

Alkalis are soluble bases that, when dissolved in a solution (typically water), produce hydroxide ions (OH⁻). These OH⁻ ions are responsible for the alkaline nature of the solution. For example, sodium hydroxide (NaOH) dissociates in water to form Na⁺ and OH⁻ ions. Alkalis are typically corrosive and feel slippery to the touch. They turn red litmus paper blue and neutralise acids by reacting with the H⁺ ions.

The Role of Water in Dissociation

Both acids and alkalis require water (or another suitable solvent) to dissociate and release their respective ions. Without a solvent, a substance like hydrogen chloride gas behaves as a covalent molecule, not an acid. Similarly, solid sodium hydroxide does not exhibit alkaline properties until it is dissolved in water, allowing the hydroxide ions to become mobile and reactive.

pH Scale and What It Shows

The pH scale runs from 0 to 14. A pH of 7 indicates a neutral solution, like pure water. Acids have pH values less than 7, and bases (alkalis) have values greater than 7.

Classifying Solutions by pH

You can classify any solution by its pH. If the value is below 7, the solution is acidic. If it's above 7, it's alkaline. Exactly 7 means the solution is neutral.

Electrolysis Fundamentals

Electrolysis is the process of using electrical energy to break down ionic compounds. For electrolysis to occur, the ionic compound must be molten or dissolved to allow its ions to be mobile. A direct current power supply drives the process, creating a positively charged anode and a negatively charged cathode.

Cation Movement and Reaction

Cations are positively charged ions. During electrolysis, cations are attracted to the negatively charged electrode, which is called the cathode. At the cathode, cations gain electrons; this process is known as reduction. For example, metal cations are often reduced to form neutral metal atoms.

Anion Movement and Reaction

Anions are negatively charged ions. During electrolysis, anions are attracted to the positively charged electrode, which is called the anode. At the anode, anions lose electrons; this process is known as oxidation. For example, halide anions are oxidised to form halogen molecules.

Summary of Ion Movement

The movement of charged ions towards oppositely charged electrodes is crucial for electrolysis. Cations (positive ions) move to the cathode (negative electrode) to be reduced, while anions (negative ions) move to the anode (positive electrode) to be oxidised. This sustained movement and reaction facilitates the decomposition of the electrolyte.