Science (Double Award) | Physical chemistry

Welcome to Physical Chemistry: The How, When, and How Fast of Reactions!

Hello future chemist! Physical chemistry might sound intimidating, but it’s really about understanding the mechanics of chemical change: How fast reactions happen (Rates), What energy changes are involved (Energetics), and Can reactions go backwards (Reversible Reactions and Equilibrium).

This chapter is essential because it explains why industry chooses specific conditions for making useful products (like fertilisers). Don't worry if this seems tricky at first—we will break down every complex idea into simple, manageable steps!

Section 1: Rates of Reaction (The Speed Limit)

The rate of reaction is simply how quickly reactants are used up or how quickly products are formed. It's measured by monitoring a change that happens during the reaction over time.

Measuring Reaction Rate

You can measure rate by tracking visible changes. Common methods include:

• Measuring the volume of gas produced: Using a gas syringe or collecting gas over water.
• Measuring the change in mass: If a gas is produced and allowed to escape, the flask loses mass over time.
• Measuring the time taken for a visual change: For example, the time taken for a precipitate to form and obscure a cross drawn beneath the flask (the "disappearing cross" experiment).

The Collision Theory: The Foundation of Rates

For any chemical reaction to happen, particles (atoms, ions, or molecules) must collide. But not just any collision will do!

A successful reaction requires two things:

1. The particles must collide in the correct orientation (they must hit each other in the right spot).
2. The particles must collide with sufficient energy, called the activation energy.

Analogy: Imagine trying to open a locked door. You have to put the key in the lock (correct orientation), and you must turn it hard enough (sufficient energy).

The rate of reaction is controlled by the frequency of successful collisions.

Factors Affecting the Rate of Reaction

If we want to speed up a reaction, we need to increase the number of successful collisions per second.

1. Temperature

When you increase the temperature:

• Particles gain kinetic energy and move faster. This increases the frequency of collisions.
• More importantly, a greater proportion of particles now have energy equal to or greater than the activation energy. This drastically increases the frequency of successful collisions.

2. Concentration (Solutions) or Pressure (Gases)

Increasing concentration (more particles in the same volume) or pressure (forcing the same number of particles into a smaller volume) means:

• The particles are closer together.
• This increases the frequency of collisions (more chances to bump into each other).
• Result: More successful collisions per second, so the rate increases.

3. Surface Area

This mainly applies to reactions involving a solid. If you use a powdered solid instead of a lump:

• The total area of solid exposed to the liquid/gas reactants increases massively.
• More particles are available to collide simultaneously.
• Result: Increased frequency of collisions, faster rate.

Quick Tip: Think of indigestion tablets. They fizz much faster when chewed (increased surface area) than when swallowed whole.

4. Use of a Catalyst

A catalyst is a substance that speeds up a reaction without being used up itself.

• It works by providing an alternative reaction pathway that requires a lower activation energy.
• Since less energy is needed, a much greater proportion of the colliding particles now have enough energy to react successfully.
• Crucial point: Catalysts do not increase the energy of the particles or the frequency of collisions; they simply make the existing collisions easier to succeed!

Section 2: Energetics (Energy Changes)

Every chemical reaction involves an exchange of energy, usually in the form of heat. We need to look at whether the reaction gives out heat or takes it in.

Energy in Reactions: Bond Breaking and Bond Forming

For a reaction to occur, existing chemical bonds in the reactants must be broken, and new bonds must be formed in the products.

• Bond Breaking: Always requires energy (endothermic process).
• Bond Forming: Always releases energy (exothermic process).

Exothermic Reactions (Energy Leaves!)

In an exothermic reaction, the energy released when new bonds form is greater than the energy required to break the old bonds.

• Key characteristic: Energy is transferred from the reaction mixture to the surroundings.
• Observation: The temperature of the surroundings increases (it feels hot).
• Examples: Combustion (burning), Neutralisation, Respiration.

Memory Aid: EXO sounds like 'exit' – energy exits the system.

Endothermic Reactions (Energy Enters!)

In an endothermic reaction, the energy required to break the old bonds is greater than the energy released when new bonds form.

• Key characteristic: Energy is transferred from the surroundings into the reaction mixture.
• Observation: The temperature of the surroundings decreases (it feels cold).
• Examples: Thermal decomposition (heating limestone), Photosynthesis, Instant cold packs.

Memory Aid: ENDO sounds like 'enter' – energy enters the system.

Activation Energy and Energy Profile Diagrams

The activation energy (\(E_a\)), which we met in Section 1, is the minimum energy required to start a reaction. It's the "hump" you have to get over for the reaction to proceed.

(In an exam, you should be able to sketch and label these diagrams)

Exothermic Diagram:

• Products are at a lower energy level than the reactants.
• The difference in energy between reactants and products is the net energy released (\(\Delta H\)).

Endothermic Diagram:

• Products are at a higher energy level than the reactants.
• The difference in energy between reactants and products is the net energy absorbed (\(\Delta H\)).

The Role of the Catalyst Revisited: A catalyst lowers the activation energy (\(E_a\)) on these diagrams, meaning the reaction peak is lower, making it easier for the reaction to happen.

Section 3: Reversible Reactions and Equilibrium

Not all reactions go completely from reactants to products. Many are reversible.

Reversible Reactions

A reversible reaction is one where the products can react together to form the original reactants again.

• We use a double arrow to represent this: \(A + B \rightleftharpoons C + D\)
• The reaction going from left to right is the forward reaction.
• The reaction going from right to left is the reverse reaction.

Example: The hydration of copper(II) sulfate is a classic school example:
\(\text{Hydrated CuSO}_4 \rightleftharpoons \text{Anhydrous CuSO}_4 + \text{Water}\)
(Blue solid) \(\rightleftharpoons\) (White solid) + (Water vapour)

Chemical Equilibrium

If a reversible reaction takes place in a closed system (where nothing can enter or leave), eventually it will reach a state of dynamic equilibrium.

When a reaction is at equilibrium:

1. The rate of the forward reaction is exactly equal to the rate of the reverse reaction. (This is why it is "dynamic"—reactions are still happening, just balancing each other out).
2. The concentrations of the reactants and products remain constant (they don't have to be equal, just unchanging).

Le Chatelier's Principle: Responding to Stress

This is the key rule for understanding equilibrium. It explains what happens if you try to change the conditions of a reaction at equilibrium.

Le Chatelier's Principle states: If a system at equilibrium is subjected to a change in conditions (a 'stress'), the system will shift its position of equilibrium to counteract that change.

Let's see how the system "fights back" against changes in concentration, temperature, and pressure.

1. Changing Concentration

• Stress: You increase the concentration of a reactant (e.g., A).
• Counteract: The system shifts the equilibrium forward (to the right) to use up the extra A. (More product is formed).
• Stress: You increase the concentration of a product (e.g., D).
• Counteract: The system shifts the equilibrium backward (to the left) to use up the extra D. (More reactant is formed).

2. Changing Temperature

You must know whether the forward reaction is exothermic or endothermic.

• Stress: You increase the temperature (you add heat).
• Counteract: The system shifts in the direction that absorbs heat—the endothermic direction.
• Stress: You decrease the temperature (you remove heat).
• Counteract: The system shifts in the direction that releases heat—the exothermic direction.

3. Changing Pressure (Only affects gases)

The system shifts to counteract a change in pressure by trying to reduce or increase the number of gas molecules. Count the total moles of gas on each side of the equation.

• Stress: You increase the pressure.
• Counteract: The system shifts to the side with the fewer moles of gas (to reduce the total volume/pressure).
• Stress: You decrease the pressure.
• Counteract: The system shifts to the side with the more moles of gas (to increase the total volume/pressure).

4. The Effect of a Catalyst on Equilibrium

A catalyst does not affect the position of equilibrium or the final yield. It only speeds up the rate at which equilibrium is achieved (it speeds up both the forward and reverse reactions equally).

Section 4: Industrial Applications (Rate vs. Yield)

In industry, chemists need to find the best conditions for a reaction. This often involves a balance between getting a high yield (good equilibrium position) and having a fast rate (good production speed).

The Compromise in the Haber Process

The Haber Process manufactures ammonia (\(\text{NH}_3\)), vital for making fertilisers, from nitrogen and hydrogen gases.

The reaction is:
\(\text{N}_{2(g)} + 3\text{H}_{2(g)} \rightleftharpoons 2\text{NH}_{3(g)}\) (\(\Delta H\) is negative / Exothermic)

Total Moles of Gas: Reactants = 4 moles. Products = 2 moles.

1. Optimal Pressure Conditions

• Le Chatelier says: To favour the product (ammonia), we want the equilibrium to shift to the side with the fewer moles of gas (the right side, 2 moles).
• Conclusion: A high pressure (around 200 atmospheres) is used. This gives a good yield.

2. Optimal Temperature Conditions

• Le Chatelier says: Since the forward reaction is exothermic, we want to remove heat. This means a low temperature would favour the product (ammonia) and give a high yield.
• But: A low temperature gives a very slow rate of reaction.
• Conclusion: Industry uses a compromise temperature (around \(450^\circ\text{C}\)). This temperature is high enough for a fast rate, even though the yield (percentage of ammonia formed) is slightly lower than if they used a very low temperature.

3. Use of Catalyst

• An iron catalyst is used.
• Reason: It does not change the yield, but it ensures that the reaction reaches its equilibrium position quickly, allowing the factory to produce ammonia faster and more efficiently.

Understanding Physical Chemistry allows us to design industrial processes that are both cost-effective and productive!