Introduction to Kinetics & Equilibria (Edexcel International A Level (IAL) Chemistry): Flashcards

A tangent is drawn to the curve at the point of interest. The gradient of the tangent (Δy ÷ Δx) gives the instantaneous rate of reaction at that time.

Rate of reaction = change in amount of reactants or products divided by time.

As a reaction proceeds, the concentration of reactants decreases and the concentration of products increases.

The gas must be sufficiently dense for the mass change to be detectable on a balance. Light gases such as hydrogen (Mr = 2.0) cause too small a mass change to measure accurately.

False.

Hydrogen has a very low M_r (2.0), so the mass loss is too small to detect on a balance. Only gases with a sufficiently high M_r, such as CO₂ (M_r = 44.0), are suitable.

Sodium thiosulfate reacts with HCl to produce a yellow precipitate of sulfur. A cross beneath the flask is timed until it can no longer be seen through the solution, giving a measure of reaction rate.

Quenching is the deliberate stopping of a reaction at a specific time so that a sample can be taken and analysed by titration without the reaction continuing during the analysis.

The gradient of the graph at any point shows the rate of reaction at that time. The gradient decreases over time as the reactant concentration falls and the reaction slows.

A Maxwell-Boltzmann distribution curve is a graph showing the distribution of kinetic energies among the particles in a gas sample at a given temperature.

The activation energy (E_a) marker divides the curve into two regions. Only particles with energy equal to or greater than E_a can undergo successful collisions and react.

In a sample of gas at a given temperature, a few particles have very low energy, a few have very high energy, and most have energy in between.

The curve flattens and the peak shifts to the right. This means a greater proportion of particles have energies above the activation energy.

False.

At higher temperature, the peak shifts to the right and the curve flattens. This reflects the higher average kinetic energy of the particles.

1. Particles move faster, leading to more frequent collisions.

2. A greater proportion of particles have kinetic energy equal to or greater than the activation energy, so more collisions are successful.

At higher temperatures, the proportion of particles with energy greater than the activation energy increases, so the rate of reaction increases.

There is always a small number of particles with very high kinetic energy. The curve approaches but never reaches zero, reflecting the statistical distribution of energies in the sample.

A catalyst is a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy. It is not consumed in the reaction.

A homogeneous catalyst is one that is in the same phase as the reactants, for example when both the catalyst and reactants are dissolved in solution.

A heterogeneous catalyst is one that is in a different phase from the reactants, for example a solid catalyst used with gaseous reactants. This is the most common type in industry.

A catalyst provides an alternative reaction pathway with a lower activation energy. A greater proportion of particles now have energy equal to or above E_a, so more collisions are successful.

False.

A heterogeneous catalyst is in a different phase from the reactants. For example, the catalyst may be a solid while the reactants are gases.

On a Maxwell-Boltzmann curve, a catalyst shifts the activation energy to a lower value, so a greater proportion of particles have energy above the activation energy.

1. They increase the rate of reaction, so more product is made in a given time.

2. They reduce energy costs as the reaction can occur at lower temperatures and pressures.

3. They can increase atom economy by reducing unwanted side reactions.

A catalyst lowers the activation energy by providing an alternative pathway. More particles in the sample have energies exceeding this lower E_a, leading to more effective collisions per unit time.

A heterogeneous catalyst is the most common type in industry, where the catalyst is in a different phase from the reactants.

A reversible reaction is one in which the products can react to reform the original reactants. It is represented using two opposing half-arrows (⇌).

Dynamic equilibrium exists in a closed system when the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of all reactants and products remain constant.

A closed system is one in which none of the reactants or products can escape from the reaction mixture. Dynamic equilibrium can only be established in a closed system.

1. The reaction must occur in a closed system.

2. The rate of the forward reaction must equal the rate of the backward reaction.

False.

At dynamic equilibrium, the concentrations of reactants and products are constant, not equal. The actual concentrations of each species can be very different from one another.

At dynamic equilibrium, the rate of the forward reaction equals the rate of the backward reaction.

In an open system, gases can escape to the surroundings. This prevents the backward reaction from occurring and causes the reaction to go to completion rather than reaching equilibrium.

False.

In a dynamic equilibrium, both the forward and backward reactions continue at the same rate. The term "dynamic" reflects that the particles are constantly reacting.

In a closed system, no reactants or products can escape the reaction mixture. In an open system, matter and energy can be lost to the surroundings, preventing equilibrium from being established.

Le Chatelier's principle states that if a change is made to a system at dynamic equilibrium, the position of equilibrium shifts to counteract that change.

The position of equilibrium refers to the relative amounts of reactants and products in an equilibrium mixture. It can shift left (towards reactants) or right (towards products).

The equilibrium shifts to the right to use up the excess reactant. This increases the concentration of products and restores balance.

Increasing pressure shifts the equilibrium towards the side with fewer moles of gas. Decreasing pressure shifts it towards the side with more moles of gas.

There is no effect on the equilibrium position. The shift in equilibrium caused by a pressure change only occurs when there are different numbers of moles of gas on each side.

False.

Increasing temperature shifts the equilibrium towards the endothermic direction. The system absorbs the extra heat energy to oppose the change.

The equilibrium shifts to the left (towards reactants). The system opposes the temperature increase by absorbing energy, which means the endothermic (reverse) reaction is favoured.

When the concentration of a product is decreased, the equilibrium shifts to the right to replace the lost product.

Removing a product decreases its concentration. By Le Chatelier's principle, the equilibrium shifts to the right to replace it, consuming more reactants and producing more product overall.

The equilibrium shifts to the left towards the side with fewer moles of gas. This reduces the pressure and opposes the change, in line with Le Chatelier's principle.

The yield-rate compromise is the balance between choosing conditions that give a high yield and conditions that give an acceptable rate of reaction. Conditions that maximise yield often reduce rate, and vice versa.

450°C gives an acceptable yield of around 35% ammonia and an acceptable rate of reaction. Lower temperatures give a higher yield but the rate is too slow; higher temperatures give a faster rate but greatly reduce the yield.

The forward reaction is exothermic, so increasing temperature shifts the equilibrium to the left, reducing the yield of ammonia. The higher energy costs also make it less economical.

False.

Increasing pressure increases the yield because there are 4 moles of gas on the reactant side but only 2 moles on the product side. The equilibrium shifts to the right towards fewer moles of gas.

In the Haber Process, there are 4 moles of gas on the reactant side and 2 moles on the product side, so increasing pressure shifts the equilibrium to the right.

Doubling the pressure to around 400 atm only increases the yield by around 7%, but greatly increases the financial and energy costs, and raises significant health and safety concerns.

1. Financial: is the cost of altered conditions balanced by sufficient increase in profit?

2. Energy and environmental: will conditions increase fossil fuel use or energy costs?

3. Health and safety: are the proposed conditions safe for workers and equipment?

In the Haber Process, the temperature compromise of approximately 450°C gives a yield of around 35 % ammonia.

At lower temperatures, the rate of reaction is too slow. Although the equilibrium yield would be higher, the time taken to produce a commercially useful quantity of ammonia would be too long.