Lesson Notes By Weeks and Term v3 - Senior Secondary 2
Acid-Base Reactions
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Acid-Base Reactions
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Subject: Chemistry
Class: Senior Secondary 2
Term: 1st Term
Week: 2
Theme: The Chemical World
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define concentration, in mol dm-3 of solutions. Define standardsolutions. Explain relationshipbetween concentrationsand volume of reactingsubstances mathematically expressthe relationship betweenthe concentration in moldm-3 and volume of as olution carry out acid-basetitrations usingappropriate in dicators record correctly titrevalues to two decimalplaces carry out relevantcalculations from titrationresults.
Concentration is a measure of the amount of solute present in a given volume of solution. When expressed in mol dm−3 (moles per cubic decimetre), it is often referred to as molarity (M).
Definition: Molarity (C or M) is defined as the number of moles of solute dissolved in one cubic decimetre (or one litre) of solution.
Units: The SI unit for volume is m3, but dm3 (litre) is commonly used in chemistry for convenience. 1 dm3 = 1000 cm3 = 1 litre.
Formula: `Molarity (C) = Number of moles of solute (n) / Volume of solution (V in dm3)` `C = n / V` Derivation of moles: `Number of moles (n) = Mass of solute (m) / Molar mass of solute (Mr)` `n = m / Mr` Combining these: `C = (m / Mr) / V` or `C = m / (Mr × V)` Worked Example 1: Calculating Molarity A student dissolved 4.9 g of tetraoxosulphate(VI) acid (H2SO4) in water to make 250 cm3 of solution. Calculate the concentration of the acid in mol dm−3. (Given: H=1, S=32, O=16)
Solution: Calculate Molar Mass (Mr) of H2SO4: H2SO4 = (2 × 1) + (1 × 32) + (4 × 16) = 2 + 32 + 64 = 98 g/mol Convert Volume (V) to dm3: V = 250 cm3 / 1000 cm3/dm3 = 0.250 dm3 Calculate Number of Moles (n) of H2SO4: n = mass / Mr = 4.9 g / 98 g/mol = 0.05 mol Calculate Concentration (C): C = n / V = 0.05 mol / 0.250 dm3 = 0.2 mol dm−3 Worked Example 2: Calculating Mass needed for a given Molarity How many grams of sodium carbonate (Na2CO3) are required to prepare 500 cm3 of a 0.1 mol dm−3 solution? (Given: Na=23, C=12, O=16)
Solution: Calculate Molar Mass (Mr) of Na2CO3: Na2CO3 = (2 × 23) + (1 × 12) + (3 × 16) = 46 + 12 + 48 = 106 g/mol Convert Volume (V) to dm3: V = 500 cm3 / 1000 cm3/dm3 = 0.5 dm3 Calculate Number of Moles (n): From C = n/V, then n = C × V n = 0.1 mol dm−3 × 0.5 dm3 = 0.05 mol Calculate Mass (m): From n = m/Mr, then m = n × Mr m = 0.05 mol × 106 g/mol = 5.3 g Definition: A standard solution is a solution whose concentration is accurately known.
Preparation: Standard solutions are prepared by dissolving an accurately weighed mass of a pure substance (called a primary standard) in a precisely known volume of solvent (usually distilled water).
Primary Standards: Substances used to prepare standard solutions directly. They must possess the following characteristics: High purity: Must be available in a pure state (e.g., >99.9%).
Stable: Must not decompose or react with air (e.g., absorb moisture (hygroscopic) or carbon dioxide) during weighing or storage.
High molar mass: To minimise errors during weighing (larger mass, smaller percentage error). Readily soluble in the solvent. Non-toxic.
Not hygroscopic or efflorescent: It should not absorb moisture from the atmosphere or lose water of crystallisation to the atmosphere.
Examples of Primary Standards: Anhydrous sodium carbonate (Na2CO3), hydrated oxalic acid (H2C2O4·2H2O), potassium hydrogen phthalate.
Secondary Standards: Solutions whose concentrations are determined by reacting them with a primary standard. Substances like NaOH (hygroscopic) and HCl (volatile) are not primary standards because they are difficult to obtain in a pure, stable form or their concentrations change easily. Their concentrations must be standardised against a primary standard. Acid-base reactions are typically stoichiometric, meaning they react in fixed molar ratios. In titration, a known volume of one solution (acid or base) reacts completely with a known volume of another solution (base or acid) of known or unknown concentration.
Principle: At the equivalence point (or stoichiometric point) of a titration, the moles of acid stoichiometrically react with the moles of base as dictated by the balanced chemical equation. `n_acid (moles of acid) = n_base (moles of base)` (if mole ratio is 1:1)
From the balanced chemical equation: `aA + bB → Products` Where 'a' and 'b' are the stoichiometric coefficients for reactants A and
B. At equivalence point: `n_A / a = n_B / b` Using C = n/V, we have n = C ×
V. Substituting this into the ratio: `(C_A × V_A) / a = (C_B × V_B) / b` Where: `C_A` = concentration of acid `V_A` = volume of acid `a` = stoichiometric coefficient of acid from balanced equation `C_B` = concentration of base `V_B` = volume of base `b` = stoichiometric coefficient of base from balanced equation This formula is crucial for calculating an unknown concentration from titration results.
Definition: Titration is a quantitative analytical method used to determine the concentration of an unknown solution by reacting it with a solution of known concentration (a standard solution). In acid-base titration, an acid reacts with a base.
Equivalence Point: The point in a titration where the moles of acid exactly neutralise the moles of base according to the stoichiometry of the reaction.
Endpoint: The point where the indicator changes colour, signaling the completion of the reaction. For a well-chosen indicator, the endpoint should be very close to the equivalence point.
Apparatus: Burette: For dispensing variable, accurate volumes of the titrant (solution of known concentration, usually placed in the burette). Reads to two decimal places.
Pipette: For accurately measuring and transferring a fixed volume of the analyte (solution of unknown concentration, usually placed in the conical flask).
Conical Flask (Erlenmeyer Flask): To hold the analyte and indicator during titration.
Retort Stand and Clamp: To hold the burette vertically.
Funnel: For filling the burette.
Wash Bottle (with distilled water): For rinsing apparatus.
White Tile/Paper: Placed under the conical flask to observe colour change clearly.
Indicator: A substance that changes colour at or near the equivalence point.
Methyl orange: Red in acidic medium, orange at neutral, yellow in alkaline medium. Used for strong acid-weak base titrations.
Phenolphthalein: Colourless in acidic medium, pink in alkaline medium. Used for strong acid-strong base or weak acid-strong base titrations.
Litmus: Red in acid, blue in base. Less precise.
General Procedure for Acid-Base Titration:
1. Preparation of Standard Solution: If not provided, accurately prepare the standard solution (e.g., 0.1 M Na2CO3) by dissolving a known mass in a volumetric flask.
2. Cleaning Apparatus: Wash the burette, pipette, and conical flask with distilled water.
3. Rinsing: Rinse the burette with a small amount of the solution to be put in it (titrant, e.g., acid). Rinse the pipette with a small amount of the solution to be measured by it (analyte, e.g., base). Rinse the conical flask with distilled water only (never with acid or base, as it changes the moles).
4. Filling the Burette: Clamp the burette vertically. Close the tap. Using a funnel, fill the burette with the titrant (acid) slightly above the zero mark. Open the tap briefly to expel air bubbles from the jet and ensure the meniscus is at or below the zero mark. Record the initial burette reading (initial titre) to two decimal places.
5. Pipetting: Use the pipette to transfer an accurate, fixed volume (e.g., 20.00 cm3 or 25.00 cm3) of the analyte (base) into the clean conical flask.
6. Adding Indicator: Add 2-3 drops of the appropriate indicator (e.g., phenolphthalein or methyl orange) to the solution in the conical flask. Note the initial colour.
7. Titration: Place the conical flask on a white tile under the burette. Slowly add the titrant from the burette to the analyte in the conical flask, swirling constantly. The titrant should be added drop-wise as the endpoint is approached (indicated by a persistent colour change that fades slowly, then becomes permanent).
8. Recording Titre Value: Stop when a single drop causes a permanent colour change (the endpoint). Record the final burette reading (final titre) to two decimal places. The difference between the initial and final readings is the titre value (volume of titrant used).
9. Repeat: Repeat the titration experiment at least two more times to obtain concordant results (titre values within ±0.1 cm3 of each other).
1
0. Calculate Average Titre: Calculate the average of the concordant titre values.
Recording Titre Values: | Titration | Initial Burette Reading (cm3) | Final Burette Reading (cm3) | Volume of Titrant (cm3) | | :-------- | :---------------------------- | :-------------------------- | :---------------------- | | Rough | 0.00 | 20.50 | 20.50 | | 1 | 0.00 | 20.25 | 20.25 | | 2 | 20.25 | 40.50 | 20.25 | | 3 | 0.00 | 20.20 | 20.20 | Average Titre = (20.25 + 20.25 + 20.20) / 3 = 20.23 cm3 (if 20.20 and 20.25 are concordant, typically within 0.1cm3 error, then
Agriculture and Soil Science (Nigeria): In many parts of Nigeria, soil acidity or alkalinity affects crop yields. Farmers use lime (calcium hydroxide) to neutralise acidic soils or ammonium sulphate to acidify alkaline soils. Titration is a key analytical technique used by agricultural scientists to determine the pH and acidity/alkalinity levels of soil samples, guiding farmers on appropriate soil amendments. For instance, soil samples from the Niger Delta region might be found to be acidic, requiring liming.
Food and Beverage Industry (Nigeria): The acidity of many locally produced food items and drinks (e.g., Zobo drink, palm wine, local yogurts, vinegars) is a crucial factor for taste, preservation, and quality control. Titration is routinely used in Nigerian food processing companies to monitor and adjust the acidity of their products to meet health standards and ensure consumer safety. For example, determining the acetic acid content in locally brewed vinegar for sale. Water Treatment and Environmental Monitoring: Water treatment plants across Nigeria use acid-base titration to monitor the pH of raw water sources and treated potable water. Maintaining a neutral pH is essential for public health and preventing corrosion of water pipes. Environmental agencies also use titration to analyse industrial effluents and ensure they meet environmental discharge standards before being released into rivers like the Niger or Benue, preventing pollution. ---