OCR GCSE Combined Science A Gateway revision notes
Chemical reactions Revision Notes
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Chemical reactions
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Understanding Dilute/Concentrated vs. Weak/Strong Acids (Higher) Revision Notes
Dilute vs. Concentrated Acids (Higher)
These terms describe the quantity of acid solute dissolved in a specific volume of solvent, typically water. A dilute acid refers to a solution where a small amount of acid molecules are present per unit volume of solution. This means the overall concentration of the acid is low. Conversely, a concentrated acid contains a large amount of acid molecules per unit volume of solution, indicating a high overall concentration. For example, 1 moldm⁻³ hydrochloric acid is more concentrated than 0.1 moldm⁻³ hydrochloric acid. Note that these terms do not describe the acid's chemical properties or reactivity in terms of ionisation.
Weak vs. Strong Acids (Higher)
These terms refer to the degree of ionisation (or dissociation) an acid undergoes when dissolved in water. Ionisation is the process where acid molecules split to produce hydrogen ions (H⁺). A strong acid is one that completely ionises (dissociates) in water. This means virtually all of its molecules break apart to release H⁺ ions. Examples include hydrochloric acid (HCl) and sulfuric acid (H₂SO₄). As a result, a strong acid solution will have a very high concentration of H⁺ ions. A weak acid is one that only partially ionises (dissociates) in water. Only a small fraction of its molecules release H⁺ ions, while most remain as complete molecules. Examples are ethanoic acid (CH₃COOH) and citric acid. Consequently, solutions of weak acids have a much lower concentration of H⁺ ions compared to strong acids of the same molar concentration.
Independence of Concentration and Strength (Higher)
It is critical to understand that 'dilute/concentrated' and 'weak/strong' are independent concepts. An acid's strength (its degree of ionisation) is an intrinsic chemical property of the acid itself. Its concentration (dilute or concentrated) is a measure of how much of that acid has been dissolved in a solvent. Therefore, you can have a dilute strong acid (e.g., a small amount of hydrochloric acid dissolved in a large volume of water; it still fully ionises) or a concentrated weak acid (e.g., a large amount of ethanoic acid dissolved in a small volume of water; it still only partially ionises). The pH of an acid solution is primarily determined by the concentration of H⁺ ions, which depends on both the acid's strength and its concentration.
Measuring pH: Indicators and pH Meters Revision Notes
Understanding pH and its measurement
The pH scale ranges from 0 to 14. A pH of 7 is neutral. Values below 7 indicate acidity, with lower numbers representing stronger acids. Values above 7 indicate alkalinity (basicity), with higher numbers representing stronger alkalis. Measuring pH is crucial in many scientific and industrial applications, from ensuring safe drinking water to controlling chemical reactions. Choosing the right measurement technique depends on the required precision.
Using pH Indicators
pH indicators are chemicals that change colour at different pH values. Universal indicator is a common mixture that displays a spectrum of colours (e.g., red for strong acid, orange/yellow for weak acid, green for neutral, blue for weak alkali, purple for strong alkali), allowing for an approximate pH determination. Litmus paper is simpler; it turns red in acid and blue in alkali, only indicating whether a solution is acidic or alkaline without giving a specific pH number. Indicators are quick and inexpensive but less precise.
Using a pH Meter
A pH meter is an electronic device that provides a precise, numerical pH reading. It consists of a glass electrode (probe) connected to a meter. To use it, the electrode is first calibrated using buffer solutions of known pH (e.g., pH 4, 7, and 10), then rinsed with distilled water, and finally immersed in the solution to be tested. The pH value is displayed digitally. pH meters offer higher accuracy and precision compared to indicators, making them suitable for situations requiring exact pH measurements.
Calculating Energy Changes in Chemical Reactions (Higher Tier) Revision Notes
Bond Breaking and Bond Forming
Chemical reactions always involve changes to chemical bonds. Existing bonds in reactant molecules must first be broken. This process always requires an input of energy from the surroundings, meaning bond breaking is an endothermic process. Following bond breaking, new bonds are formed to create the product molecules. The formation of new bonds always releases energy to the surroundings, making bond formation an exothermic process. The overall energy change of a reaction is the net difference between these two energy transfers.
Calculating Overall Energy Change (Higher Tier)
To determine the overall energy change (ΔH) for a chemical reaction, you need to use average bond energy values. The calculation follows this formula: ΔH = Σ(Bond energies of bonds broken in reactants) - Σ(Bond energies of bonds formed in products). If the calculated ΔH is positive, the reaction is endothermic, meaning more energy was absorbed to break bonds than released during bond formation. If ΔH is negative, the reaction is exothermic, indicating that more energy was released during bond formation than was absorbed to break bonds.
Exothermic vs. Endothermic Reactions
An exothermic reaction is one where energy is released to the surroundings, causing the temperature of the surroundings to increase. In terms of bond energies, this occurs when the energy released from forming new bonds is greater than the energy absorbed to break existing bonds. The enthalpy change (ΔH) for an exothermic reaction is negative. Conversely, an endothermic reaction is one where energy is absorbed from the surroundings, causing the temperature of the surroundings to decrease. This happens when the energy absorbed to break existing bonds is greater than the energy released from forming new bonds. The enthalpy change (ΔH) for an endothermic reaction is positive.
Understanding pH, Acidity, Alkalinity, and Neutrality (Higher Tier) Revision Notes
The pH Scale and Hydrogen Ion Concentration (Higher Tier)
The pH scale is a numerical measure of acidity or alkalinity, typically ranging from 0 to 14. It is directly linked to the concentration of hydrogen ions (H⁺) in a solution. A lower pH value indicates a higher concentration of H⁺ ions, signifying an acidic solution. Conversely, a higher pH value indicates a lower concentration of H⁺ ions (and a higher concentration of hydroxide ions, OH⁻), signifying an alkaline solution. Neutral solutions have a pH of 7, with equal concentrations of H⁺ and OH⁻ ions.
Relative Acidity and Alkalinity (Higher Tier)
A solution with a pH below 7 is acidic, and the further the pH is below 7, the more acidic it is. For example, pH 1 is strongly acidic, while pH 6 is weakly acidic. A solution with a pH above 7 is alkaline, and the further the pH is above 7, the more alkaline it is. For instance, pH 14 is strongly alkaline, while pH 8 is weakly alkaline. The relationship between pH and hydrogen ion concentration is logarithmic; a decrease of one pH unit means a tenfold increase in H⁺ concentration.
Neutrality (Higher Tier)
A neutral solution has a pH of exactly 7. At this point, the concentration of hydrogen ions (H⁺) is equal to the concentration of hydroxide ions (OH⁻). Pure water at 25°C is a common example of a neutral substance. Reactions between strong acids and strong alkalis typically produce neutral solutions when completely neutralised, although the resulting salt solution may not always be pH 7 due to hydrolysis of ions.
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