Chemical reactions Revision Notes

What are Acids?

Acids are chemical substances that, when dissolved in water, produce hydrogen ions (H⁺). These hydrogen ions are fundamental to the definition and behaviour of acids. For example, hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃) all release H⁺ ions when they come into contact with water. The presence of these H⁺ ions is what gives acidic solutions their characteristic properties, such as a pH value less than 7 and the ability to turn blue litmus paper red. It's important to remember that the more H⁺ ions an acid produces in solution, the stronger the acid and the lower its pH will be.

What are Alkalis?

Alkalis are soluble bases that release hydroxide ions (OH⁻) when dissolved in water. Common examples include sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide (Ca(OH)₂). The hydroxide ions are responsible for the typical properties of alkaline solutions, such as a pH value greater than 7 and the ability to turn red litmus paper blue. Alkalis are often described as having a 'slippery' feel. The concentration of OH⁻ ions is directly related to the alkalinity of the solution; a higher concentration means a stronger alkali and a higher pH value.

Key Ions in Acid-Base Chemistry

The behavior of acids and alkalis is primarily determined by their characteristic ions. Acids always involve the production of hydrogen ions (H⁺) in aqueous solutions. These H⁺ ions are a proton and are highly reactive. Alkalis always involve the production of hydroxide ions (OH⁻) in aqueous solutions. These ions are a combination of an oxygen atom and a hydrogen atom, carrying a single negative charge. The balance between H⁺ and OH⁻ concentrations in a solution dictates its pH and whether it is acidic, neutral, or alkaline. In neutral solutions, the concentrations of H⁺ and OH⁻ are equal.

What is an Ionic Equation? (Higher Tier)

An ionic equation shows only the particles (ions or compounds) that participate chemically in a reaction, excluding spectator ions. Spectator ions are present in the reaction mixture but do not take part in the actual chemical change; they remain unchanged from reactants to products. Ionic equations are particularly useful for understanding reactions involving electrolytes in solution, such as precipitation, acid-base, and redox reactions. This simplification helps to highlight the essential chemical transformation occurring.

Steps to Construct a Balanced Ionic Equation (Higher Tier)

1. Write a full balanced chemical equation for the reaction, including state symbols (s, l, g, aq).
2. Dissociate all aqueous ionic compounds into their separate ions. Remember to keep charges correct and balance the number of each ion. Do not dissociate solids, liquids, gases, or covalent compounds.
3. Identify and cancel out spectator ions. Spectator ions are identical on both sides of the equation.
4. Write the remaining species as the net ionic equation. This equation must show both atom and charge balance. For example, if you have two positive charges on the left, you must have two positive charges on the right.

Identifying Ions and Charges (Higher Tier)

To correctly dissociate compounds, you need to know common ion charges. Group 1 metals typically form +1 ions (e.g., Na+), Group 2 metals form +2 ions (e.g., Mg2+), and Group 17 non-metals (halogens) form -1 ions (e.g., Cl-). Polyatomic ions like sulfate (SO4^2-) and nitrate (NO3^-) also have specific charges. For example, sodium chloride (NaCl) is Na+ and Cl-, while magnesium sulfate (MgSO4) is Mg2+ and SO4^2-. Remember that the overall charge of an ionic compound is zero, which helps in deducing individual ion charges if you know one.

Balancing the Ionic Equation (Higher Tier)

The final ionic equation must be balanced in two ways: by atoms and by charge. Atom balance means the number of atoms of each element must be the same on both sides of the equation. Charge balance means the sum of the charges on the reactant side must equal the sum of the charges on the product side. For example, if you have 2Na+(aq) on the left, and Na+ is a spectator, you must have 2Na+(aq) on the right. If the net charge on the left is +1, the net charge on the right must also be +1. This ensures that the equation accurately represents the conservation of mass and charge during the reaction.

What is Neutralisation?

Neutralisation is a chemical reaction that occurs when an acid reacts with an alkali (a soluble base) or any base. The key products of this reaction are a salt and water. This process effectively cancels out, or 'neutralises', the corrosive properties of both the acid and the base, leading to a solution that is typically closer to neutral pH. For instance, when an acid like hydrochloric acid (HCl) mixes with an alkali like sodium hydroxide (NaOH), the resulting products are sodium chloride (NaCl), which is a salt, and water (H2O).

Reactants and Products

The general word equation for a neutralisation reaction is: Acid + Alkali/Base → Salt + Water. The acid donates hydrogen ions (H+), and the alkali or base accepts these ions, often providing hydroxide ions (OH-) to combine with H+ to form water. The remaining positive ion from the base/alkali and the negative ion from the acid combine to form a salt. For example, nitric acid + potassium hydroxide → potassium nitrate + water. The type of salt formed depends entirely on the specific acid and base/alkali used as reactants.

Example: Hydrochloric Acid and Sodium Hydroxide

A common example of neutralisation is the reaction between hydrochloric acid and sodium hydroxide. Hydrochloric acid (HCl) is a strong acid, and sodium hydroxide (NaOH) is a strong alkali. When they react, the hydrogen ions from HCl combine with the hydroxide ions from NaOH to form water (H2O). The sodium ions (Na+) and chloride ions (Cl-) then combine to form sodium chloride (NaCl), which is a common table salt. The word equation for this specific reaction is: Hydrochloric acid + Sodium hydroxide → Sodium chloride + Water.

The Role of Water in Neutralisation

In neutralisation reactions, the formation of water is a crucial aspect. Acids produce hydrogen ions (H+), and alkalis produce hydroxide ions (OH-). These two ions readily react to form water molecules (H2O). This removal of both H+ and OH- ions from the solution is what causes the pH to shift towards neutral. The other product, the salt, is formed from the spectator ions (the ions that do not participate directly in the formation of water) from the acid and the base.

Identifying Carbon Dioxide (CO₂)

To test for carbon dioxide, you should bubble the gas through limewater, which is a solution of calcium hydroxide. If carbon dioxide is present, the limewater will turn cloudy or milky. This happens because carbon dioxide reacts with calcium hydroxide to form insoluble calcium carbonate, which is the white precipitate observed. The chemical reaction is: CO₂(g) + Ca(OH)₂(aq) → CaCO₃(s) + H₂O(l).

Identifying Hydrogen (H₂)

Hydrogen gas is highly flammable. To test for hydrogen, you should hold a lit splint at the mouth of a test tube containing the gas. If hydrogen is present, the lit splint will be extinguished with a characteristic 'squeaky pop' sound. This distinctive sound is due to the rapid combustion of hydrogen in the presence of oxygen, forming water. The chemical reaction is: 2H₂(g) + O₂(g) → 2H₂O(g).

Identifying Oxygen (O₂)

Oxygen gas supports combustion but does not burn itself. To test for oxygen, you should insert a glowing splint into a test tube containing the gas. If oxygen is present, the glowing splint will relight. This observation is due to oxygen's ability to significantly increase the rate of combustion of the wood in the splint. No specific chemical equation is usually given for this observation, as it's a physical observation of supported combustion.

Identifying Chlorine (Cl₂)

Chlorine gas is a greenish-yellow gas with a pungent smell, known for its bleaching properties. To test for chlorine, you should hold a piece of damp blue litmus paper in the gas. If chlorine is present, the blue litmus paper will turn red and then be bleached white. The initial red colour is due to chlorine dissolving in water to form hydrochloric acid and chloric(I) acid, both acidic. The subsequent bleaching is due to the chloric(I) acid, a strong oxidising agent. The reaction is: Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq).