Rates of Chemical Reactions 🌡️

Definition: Rate of chemical reaction is the speed at which the reaction takes place. ⏱️

Formula: Rate of reaction = (change in concentration of reactant or product) / time 📊

Measuring the Rate of Reaction 📏

• How quickly a product is obtained. 🚀
• How quickly a reactant is used up. ⏳

Example 📚

Consider the reaction below:

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

The rate of reaction can be measured by:

• The volume of carbon dioxide over time. 💨
• The decrease in mass of the system due to loss of carbon dioxide. ⚖️

[A] Measuring the Rate by Volume of Gas Produced 💨

A graduated syringe is used to measure the volume of carbon dioxide gas formed over time. 📏

The total volume of carbon dioxide given off at one-minute intervals is recorded, and a graph of total volume of carbon dioxide against time is plotted. The gradient of this graph represents the rate of reaction. 📈

[B] Measuring the Rate by Decrease in Mass ⚖️

A mass balance is used to follow the loss in mass as carbon dioxide escapes. Mass readings are taken at one-minute intervals and plotted against time. The gradient of the graph at various points represents the reaction rate. 📉

Note: Cotton wool acts as a stopper, allowing carbon dioxide to escape while preventing splashing. 🧶

Collision Theory 💥

The Collision Theory states that particles must collide with a certain amount of energy for a reaction to occur. ⚡

Factors Affecting the Rate of Reaction ⚙️

• Concentration: Increasing concentration raises reaction rate due to more frequent collisions. 🔬
• Temperature: Higher temperatures increase kinetic energy, causing more effective collisions. 🔥
• Pressure: In gaseous reactions, increased pressure forces particles closer, increasing collisions. 💨
• Surface Area: Smaller particle sizes offer greater surface area, enhancing reaction rate. 🪨
• Catalyst: Catalysts increase the reaction rate without being consumed. ⚗️

Example Data Table 📊

Bond Energy 🔗

Bond energy is the energy required to break a bond or the energy released when a bond forms. The enthalpy change (ΔH) of a reaction depends on bond energies in reactants and products. ⚡

Formula: ΔH = Σ(bond energies of reactants) - Σ(bond energies of products) 📏

Example Calculation of Bond Energy 🧮

ΔH = (436 + 243) - (2 x 431) = 679 - 862 = -183 kJ/mol 🔍

Equilibrium is the state where forward and reverse reaction rates are equal, and concentrations remain constant in a closed system. 🔄

Le Chatelier’s Principle states that if a system at equilibrium is disturbed, it will shift in a direction that counteracts the disturbance. This principle applies to changes in:

• Temperature 🌡️: Increasing temperature favors the forward reaction of an endothermic process, while a decrease in temperature favors the reverse (exothermic) reaction.
• Concentration 🔬: Increasing the concentration of reactants favors the forward reaction, while increasing the concentration of products favors the reverse reaction.
• Pressure 💨: In reactions involving gases, an increase in pressure favors the side with fewer gas molecules, reducing volume.

Example of Equilibrium Shift in the Haber Process ⚗️

The reaction of nitrogen and hydrogen to form ammonia (Haber process) is sensitive to pressure changes:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

• An increase in pressure will shift the equilibrium toward producing more ammonia, as it reduces the volume. 📈
• Decreasing pressure favors the reverse reaction, producing nitrogen and hydrogen. 📉

Symbol: Ea. Activation energy is the minimum energy needed to initiate a reaction. 🔋

Effects of Catalysts on Equilibrium ⚙️

A catalyst speeds up both the forward and reverse reactions equally, allowing equilibrium to be achieved more quickly. ⏩

Characteristics of Catalysts 🌟

• Catalysts affect the rate but not the equilibrium position. ⚖️
• They remain chemically unchanged after the reaction. 🔄
• Only small amounts are required to be effective. 📏
• They can be poisoned by impurities, reducing effectiveness. ⚠️

Examples of Industrial Catalysts 🏭

Energetics refers to the study of energy changes in chemical reactions. 🔋

Enthalpy (H) 🔥

The total energy contained in one mole of a substance, often called the heat content. 🔥

Types of Energy Changes 🔄

Endothermic Reactions ❄️

• Definition: Reactions that absorb heat energy from the surroundings, causing the temperature to fall. 🌡️
• Examples: Photosynthesis, dissolving ammonium nitrate in water. 🌱
• In an endothermic reaction, ΔH is positive because reactants have less energy than products. ➕

Exothermic Reactions 🔥

• Definition: Reactions that release heat energy to the surroundings, causing the temperature to rise. 🔥
• Examples: Combustion of fuels, respiration. 🔥
• In an exothermic reaction, ΔH is negative because reactants have more energy than products. ➖

Example Calculation 🧮

Consider the reaction between hydrogen and chlorine gases to form hydrogen chloride (HCl):

H₂ + Cl₂ → 2HCl

Using the following bond energies, we can calculate ΔH:

ΔH = (436 + 243) - (2 x 431) = 679 - 862 = -183 kJ/mol

Exercises 📝

• Question 1: Explain why increasing the surface area of a reactant increases the reaction rate. 📈
• Question 2: Describe the effect of a catalyst on the activation energy of a reaction. ⚗️
• Question 3: Calculate the enthalpy change for a reaction given the specific bond energies of reactants and products. 📊

Summary 📚

This document has explored rates of chemical reactions, the Collision Theory, factors affecting reaction rate, and equilibrium principles. We also discussed energy changes in reactions, including endothermic and exothermic reactions, as well as bond energy calculations. 🔍