Rate of Reaction

This chapter asks you to define rate of reaction, state five factors that affect it, and give examples of its applications in daily life and industries p. 114. Toast is made through the Maillard reaction, a fast reaction in which carbohydrate reacts with protein to form Amadori compounds p. 115.

4.1 Introduction to Rate of Reaction

A chemical reaction is a process in which one or more reactants are converted to one or more products, such as colourless potassium iodide and lead(II) nitrate solutions forming yellow lead(II) iodide precipitate p. 116. During a reaction, the quantity of the reactant decreases while the quantity of the product increases p. 116. A fast reaction has a high rate and takes a shorter time to complete, as in the burning of butane gas, while a slow reaction such as rusting of iron takes a longer time p. 117.

Rate of reaction is the change in the quantity of reactant or product per unit time p. 118. It can be measured from changes in mass, volume or concentration, pressure, temperature, pH value, conductivity or colour intensity, the volume of gas released or the height of precipitate formed p. 118. When 0.3 g of magnesium tape reacts completely with excess dilute hydrochloric acid in 30 s, the rate of reaction is 0.01 g s⁻¹ p. 118.

The average rate of reaction is the average rate in a specific time interval, such as 20.00 cm³ of hydrogen gas in the first 60 s giving 0.33 cm³ s⁻¹ p. 119. For the whole reaction, 35.00 cm³ of gas collected in 180 s gives 0.19 cm³ s⁻¹, because the reaction ends at 180 s p. 120. The rate of reaction at a particular point of time equals the gradient of the tangent to the curve at that time p. 121. For zinc with dilute hydrochloric acid, the rate at the 60th second is 0.25 cm³ s⁻¹ and at the 120th second is 0.11 cm³ s⁻¹ p. 123.

4.2 Factors Affecting Rate of Reaction

The rate of reaction increases when the temperature of reactants increases, when a catalyst is used, when the concentration of reactants increases, when pressure increases for gaseous reactants, and when the size of solid reactants decreases p. 125. In Experiment 4.1, the hypothesis is that the higher the temperature of reactants, the higher the rate of reaction p. 125. Sodium thiosulphate solution of 0.2 mol dm⁻³ reacts with 1 mol dm⁻³ sulphuric acid in a conical flask placed on a white paper marked ‘X’ p. 126. The solution is repeated at 35 °C, 40 °C, 45 °C and 50 °C, and graphs of temperature against time and against 1/time are drawn p. 127.

In Experiment 4.2, sodium thiosulphate solutions of 0.20 to 0.04 mol dm⁻³ test the hypothesis that the higher the concentration of reactants, the higher the rate of reaction p. 128. In Experiment 4.3, large and small marble chips react with 0.1 mol dm⁻³ dilute hydrochloric acid and the time taken to collect 30.00 cm³ of gas is measured p. 129. Large marble chips of 2 g react with 40 cm³ of acid, and the gas is collected in an inverted burette p. 130. In Experiment 4.4, copper(II) sulphate solution acts as a catalyst for the reaction between zinc and dilute hydrochloric acid p. 131. The mixture of hydrogen and air in the burette can explode when ignited, so the gas must not be ignited p. 132. For reactions involving gaseous reactants, the rate of reaction usually increases when pressure increases p. 134.

4.3 Applications of the Concept of Rate of Reaction

In daily life and industries, factors that affect the rate of reaction are adjusted, for example a refrigerator lowers the temperature of food to slow down food spoilage p. 134. In the Haber Process, nitrogen and hydrogen in the ratio 1:3 at 450 °C to 550 °C and 200 atm pass over iron filings as a catalyst to produce ammonia p. 135. In the Contact Process, sulphur is burnt to sulphur dioxide, which mixes with excess air at 450 °C and 1 atm over vanadium(V) oxide to produce sulphur trioxide p. 135. Sulphur trioxide dissolves in concentrated sulphuric acid to form oleum, which is diluted with water to produce concentrated sulphuric acid p. 136.