Monitoring and controlling chemical reactions Revision Notes

Surface Area and Collisions

Chemical reactions occur when reactant particles collide with sufficient energy. For reactions involving a solid, only the particles on the surface of the solid are available to collide with particles of other reactants (e.g., liquids or gases). If a solid is broken into smaller pieces, its total surface area increases significantly. This larger surface area means more solid particles are exposed. With more exposed particles, there is a higher chance for successful collisions between the solid reactant and other reactants, leading to a faster rate of reaction. Think of it like a crowded room – more people (particles) means more chances for bumping into each other (collisions).

Surface Area to Volume Ratio

The concept of surface area to volume ratio is crucial. When a solid object is divided into smaller pieces, its total volume remains the same, but its total surface area increases substantially. For example, a single large cube has a certain surface area. If you cut that cube into eight smaller cubes, the total volume of all eight smaller cubes is still the same as the original large cube, but the combined surface area of the eight smaller cubes is double that of the original large cube. A higher surface area to volume ratio means a greater proportion of the reactant particles are exposed and available for reaction, leading to an increased reaction rate.

Examples in Reactions

Consider the reaction of a metal with an acid. If you use a large lump of metal, the reaction will be slow because only the outer surface of the metal is exposed to the acid. However, if you use the same mass of the metal in powdered form, the reaction will be much faster. This is because the powder has a vastly larger surface area exposed to the acid, leading to more frequent collisions between the acid particles and the metal particles. Another example is burning fine wood dust versus a large log – the dust burns much more rapidly due to its larger exposed surface.

What is a Reversible Reaction?

A reversible reaction is a chemical process where the reactants form products, and at the same time, the products can react together to reform the original reactants. These reactions are represented by a double arrow (⇌) in the chemical equation, indicating that the reaction proceeds in both directions.

The Concept of a Closed System

Dynamic equilibrium can only be established in a closed system. A closed system is defined as a system where no matter (reactants or products) can enter or leave. Energy, however, can be exchanged with the surroundings. This containment is essential to allow the concentrations of all species to reach a steady state, preventing their loss or dilution.

Conditions for Dynamic Equilibrium

Dynamic equilibrium is achieved when two critical conditions are met. Firstly, the rate at which the forward reaction proceeds must become equal to the rate at which the reverse reaction proceeds. Secondly, this equality of rates must occur in a closed system. When these conditions are fulfilled, the concentrations of reactants and products remain constant, even though both reactions are still actively occurring.

Characteristics of Dynamic Equilibrium

At dynamic equilibrium, the forward and reverse reactions continue to happen, but because their rates are equal, there is no observable change in the macroscopic properties of the system. This means that concentrations, pressure, and temperature remain constant over time. The system appears static on a large scale, but at a molecular level, there is continuous conversion between reactants and products.

What are Reversible Reactions?

Reversible reactions are chemical reactions where the products can react together to reform the original reactants. This is different from irreversible reactions, which proceed in only one direction. A double arrow (⇌) is used in chemical equations to show that a reaction is reversible, indicating that the forward and reverse reactions are happening simultaneously.

Examples of Reversible Reactions

A common example of a reversible reaction involves the hydration and dehydration of anhydrous copper(II) sulfate. When white anhydrous copper(II) sulfate is mixed with water, it absorbs the water to form blue hydrated copper(II) sulfate. Conversely, when blue hydrated copper(II) sulfate is heated, it loses its water, reverting to white anhydrous copper(II) sulfate. Another example is ammonium chloride decomposition and recombination: NH₄Cl(s) ⇌ NH₃(g) + HCl(g).

How to Reverse a Reaction

The direction of a reversible reaction can be controlled by changing the reaction conditions. Key conditions that can be altered include temperature, pressure (especially for reactions involving gases), and the concentration of reactants or products. For instance, heating a reversible reaction generally favours the endothermic direction, as heat is absorbed. Cooling often favours the exothermic direction, where heat is released. Removing a product or adding more of a reactant can also shift the equilibrium position, pushing the reaction in a desired direction.

Understanding Reaction Rate Graphs

Rate of reaction graphs typically plot an amount of reactant or product (on the y-axis) against time (on the x-axis). The shape of the curve provides key information about how the reaction is proceeding. If the y-axis measures product formed, the curve will rise from zero. If it measures reactant remaining, the curve will fall from an initial maximum value. A steeper gradient indicates a faster reaction rate, as more product is formed or more reactant is consumed in a given time interval. A shallow gradient means a slower reaction, and a horizontal line indicates the reaction has stopped.

Interpreting Graph Features

Several features on a reaction rate graph are crucial for interpretation. The initial slope of the curve (at time = 0) represents the initial rate of reaction, which is typically the fastest. As the reaction proceeds, the curve usually becomes less steep, showing that the reaction is slowing down because reactants are being used up. The point where the curve becomes completely flat (horizontal) signifies that the reaction has finished. At this point, no more product is being formed, and no more reactant is being consumed. The total amount of product formed or reactant consumed can be read from the final value on the y-axis.

Comparing Reaction Rates

When comparing two or more reactions shown on the same graph, look at the steepness of their curves. The reaction with the steepest initial gradient is the fastest. If multiple reactions produce different final amounts of product, their horizontal plateaus will be at different y-axis values. A higher plateau for product formed means more product was generated. For example, if you are comparing the effect of temperature, the higher temperature reaction will show a steeper curve and reach completion faster, assuming the same amount of limiting reactant is present. If you increase the concentration of the limiting reactant, the final plateau will be higher as more product will be formed.