Monitoring Chemical Reactions

Calculating the Mass of the Product

A theoretical yield is the maximum possible mass of a product that can be made in a chemical reaction.

It can be calculated using:

• the balanced equation

• the mass and relative formula mass of the limiting reactant

• the relative formula mass of the product

Practice Question:

Lithium hydroxide is used to absorb exhaled carbon dioxide in spacecraft:

2LiOH(s) + CO₂(g) → Li₂CO₃(s) + H₂O(l)

Calculate the maximum mass of water that can be made from an excess of carbon dioxide and 95.6 g of lithium hydroxide. (Relative atomic masses: H = 1.0, Li = 6.9, O = 16.0)

relative formula mass, M_r, of LiOH = 6.9 + 16.0 + 1.0 = 23.9

relative formula mass, M_r, of H₂O = (2 × 1.0) + 16.0 = 18.0

Looking at the balanced equation:

• sum of M_r for LiOH = (2 × 23.9) = 47.8

• sum of M_r for H₂O = 18.0

• = 36.0g

Percentage Yield

An actual yield is the mass of a product actually obtained from the reaction. It is usually less than the theoretical yield. The reasons for this include:

• incomplete reactions, in which some of the reactants do not react to form the product

• practical losses during the experiment, such as during pouring or filtering

• side reactions (unwanted reactions that compete with the desired reaction)

Calculating percentage yield

The percentage yield is calculated using this equation:

percentage yield = (actual yield)/(theoretical yield) * 100

Key fact

The percentage yield can vary from 100% (no product has been lost) to 0% (no product has been made)

Question:

Copper oxide reacts with sulfuric acid to make copper sulfate and water. In an experiment, 1.6 g of dry copper sulfate crystals are made. If the theoretical yield is 2.0 g, calculate the percentage yield of copper sulfate.

actual yield = 1.6 g

percentage yield = 1.6/2.0 × 100

percentage yield = 80%

Atom Economy

No atoms are created or destroyed in a chemical reaction. However, the atoms in the reactants may not become the desired product. They may instead end up forming other products, which are regarded as waste products, also called by-products.

For example, hydrogen can be manufactured by reacting methane with steam:

methane + steam → hydrogen + carbon monoxide

CH₄(g) + H₂O(g) → 3H₂(g) + CO(g)

In this reaction, carbon and oxygen atoms in the reactants do not form the useful product. Carbon monoxide is a waste gas.

The atom economy of a reaction is a measure of how many reactant atoms form a desired product.

Calculating atom economy

The atom economy of a reaction is calculated using this equation:

atom economy = (total M of the desired product)/(total M of all reactants) × 100

Key fact

The maximum atom economy possible for a reaction is 100%. This will be the case if there is only one product (the desired product) and no by-products.

Question:

Hydrogen can be manufactured by reacting methane with steam:

CH₄(g) + H₂O(g) → 3H₂(g) + CO(g)

Calculate the atom economy for the reaction. (Relative atomic masses: H = 1.0, C = 12.0, O = 16.0)

M_r of CH₄ = 12.0 + (4 × 1.0) = 16.0

M_r of H₂O = (2 × 1.0) + 16.0 = 18.0

total M_r of reactants = 16.0 + 18.0 = 34.0

A_r of H₂ = (2 × 1.0) = 2.0

total M_r of desired product = 3 × 2.0 = 6.0 (there are three H₂ in the balanced equation)

atom economy = 17.6% (to 3 significant figures)

Practical Activity

Titration

It is important in this core practical to use appropriate apparatus to make and record a range of volume measurements accurately. This includes the safe use and handling of liquids, and monitoring chemical changes.

This outlines one way to carry out the practical. Eye protection must be worn.

Aims

To carry out an accurate titration using dilute hydrochloric acid, dilute sodium hydroxide solution, and phenolphthalein indicator.

Method

1. Use a pipette and pipette filler to add 25 cm³ of dilute sodium hydroxide solution to a clean conical flask.

2. Add a few drops of phenolphthalein indicator and put the conical flask on a white tile.

3. Fill the burette with dilute hydrochloric acid and note the starting volume.

4. Slowly add the acid from the burette to the conical flask, swirling to mix.

5. Stop adding the acid when the end-point is reached (when the colour first permanently changes from pink to colourless). Note the final volume reading.

6. Repeat steps 1 to 5 until you get concordant titres (see the Analysis).

Results

Record the results in a suitable table. The one here also shows some sample readings.

Readings should be recorded to two decimal places, ending in 0 or 5 (where the liquid level is between two graduations on the burette). The titre is the volume added (the difference between the end and start readings).

Analysis

Tick (✓) at least two concordant titres. These are titres within 0.20 cm³ (or sometimes 0.10 cm³) of each other.

Converting Units

The concentration of a solution can be measured in g/dm³ or in mol/dm³.

Volume units

Volumes used in concentration calculations must be in dm³, not in cm³. It is useful to know that 1 dm³ = 1,000 cm³. This means:

1. divide by 1,000 to convert from cm³ to dm³

2. multiply by 1,000 to convert from dm³ to cm³

For example, 500 cm³ is 0.5 dm³ (500 ÷ 1,000). It is often easiest to convert from cm³ to dm³ before continuing with a concentration calculation.

Converting between units

The relative formula mass, M_r, of the solute is used to convert between mol/dm³ and g/dm³:

To convert from mol/dm³ to g/dm³, multiply by the M_r.

Example

Calculate the concentration of 0.100 mol/dm³ sodium hydroxide solution in g/dm³. (Relative formula mass: NaOH = 40.0)

concentration = 0.100 × 40.0

= 4.00 g/dm³

Controlling Chemical Reactions

Rate of Reaction

The rate of a reaction is a measure of how quickly a reactant is used up or a product is formed.

Collision theory

For a chemical reaction to happen:

• reactant particles must collide with each other

• the particles must have enough energy for them to react

A collision that produces a reaction is called a successful collision. The activation energy is the minimum amount of energy needed by particles for a collision to be successful. It is different for different reactions.

Measuring rates of reaction

There are different ways to determine the rate of a reaction. The method chosen usually depends on the reactants and products involved, and how easy it is to measure changes in them.

In addition, the rate of reaction determines how long a reaction is observed. Reactions can vary from being almost instant to taking years to complete. In the lab, reactions are usually followed over a few seconds or minutes.

Rates, Concentration, and Pressure

The greater the rate or frequency of successful collisions, the greater the rate of reaction. If the concentration of a reacting solution or the pressure of a reacting gas is increased:

• the reactant particles become more crowded

• the frequency of collisions between reactant particles increases

• the rate of reaction increases

Note that the mean energy of the particles does not change. However, since the frequency of collisions increases, the frequency of successful collisions also increases.

Graphs

The rates of two or more reactions can be compared using a graph of mass or volume of product formed against time. The graph shows this for two reactions.

Comparing reactions of different concentration and pressure

The gradient of the line is equal to the rate of reaction. The faster reaction at the higher concentration or pressure:

• gives a steeper line

• finishes sooner

Key fact

Make sure you answer questions in terms of frequency or rate of successful collisions, rather than the number of successful collisions. This is because, given enough time, even a slow reaction will have a large number of collisions.

Rates and Surface Area to Volume Ratio

Dividing lumps

For a given mass of a solid, large lumps have smaller surface area to volume ratios than smaller lumps or powders. If a large lump is divided or ground into a powder:

• its total volume stays the same

• the area of exposed surface increases

• the surface area to volume ratio increases

Lumps vs powders

The greater the rate or frequency of successful collisions, the greater the rate of reaction. If the surface area to volume ratio of a reacting solid is increased:

• more reactant particles are exposed at the surface

• the frequency of successful collisions between reactant particles increases

• the rate of reaction increases

Note that the mean energy of the particles does not change. However, since the frequency of collisions increases, the frequency of successful collisions also increases.

Graphs

The rates of two or more reactions can be compared using a graph of mass or volume of product formed against time. The graph shows this for two reactions.

Comparing reactions of different surface area

The gradient of the line is equal to the rate of reaction. The faster reaction with the powder:

• gives a steeper line

• finishes sooner

Key fact

Make sure you answer questions in terms of surface area to volume ratio, rather than just surface area. This is because the surface area also depends on the mass of solid reactant used.

Equilibria

Reversible Reactions

In principle, all chemical reactions are reversible reactions. The products could be changed back into the original reactants using a suitable reaction. This is not obvious when a reaction ‘goes to completion’, a situation in which very little or no reactants are left. Examples of reactions that go to completion are:

• complete combustion of a fuel

• many precipitation reactions

• reactions in which a product escapes, usually a gas

It is more obvious in reactions that do not go to completion that the reaction is reversible. The reaction mixture may contain reactants and products, and their proportions may be changed by altering the reaction conditions.

Two examples

Ammonium chloride

Ammonium chloride is a white solid. It breaks down when heated, forming ammonia and hydrogen chloride. When these two gases are cool enough, they react together to form ammonium chloride again. This reversible reaction can be modelled as:

ammonium chloride ⇌ ammonia + hydrogen chloride

NH₄Cl(s) ⇌ NH₃(g) + HCl(g)

The symbol ⇌ has two half arrowheads, one pointing in each direction. It is used in equations that model reversible reactions:

• the forward reaction is the one that goes to the right

• the backward reaction is the one that goes to the left

Copper sulfate

Blue copper sulfate is described as hydrated. The copper ions in its crystal lattice structure are surrounded by water molecules. This water is driven off when blue hydrated copper sulfate is heated, leaving white anhydrous copper sulfate. This reaction is reversible:

hydrated copper sulfate ⇌ anhydrous copper sulfate + water

CuSO₄.5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)

Dynamic equilibrium

When a reversible reaction happens in a closed system, such as a stoppered flask, it reaches a dynamic equilibrium. At equilibrium:

• the forward and backward reactions are still happening

• the forward and backward reactions have an equal rate of reaction

• the concentrations of all the reacting substances remain constant and do not change

Changing the Position of Equilibrium

The equilibrium position of a reversible reaction is a measure of the concentrations of the reacting substances at equilibrium. Using the Haber process, which makes ammonia, as an example:

nitrogen + hydrogen ⇌ ammonia

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

The equilibrium position is:

• to the left if the concentrations of N₂ and H₂ are greater than the concentration of NH₃

• to the right if the concentration of NH3 is greater than the concentrations of N₂ and H₂

The equilibrium position can be changed by changing the reaction conditions by:

• changing the pressure

• changing the concentration

• changing the temperature

Changing the pressure

If the pressure is increased in a reaction involving gases, the equilibrium position moves in the direction of the fewest molecules of gas.

There are fewer molecules on the right hand side of the equation for the Haber process:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

1 + 3 = 4 molecules / 2 molecules

If the pressure is increased, the equilibrium position moves to the right.

Changing the concentration

If the concentration of a solute is increased in a reaction involving solutions, the equilibrium position moves in the direction away from this solute. For example, bismuth chloride reacts with water in a reversible reaction:

BiCl₃(aq) + H₂O(l) ⇌ BiOCl(s) + 2HCl(aq)

The concentration of hydrochloric acid can be increased by adding more hydrochloric acid. When this happens, the equilibrium position moves to the left, away from HCl(aq) in the equation.

Changing the temperature

In a reversible reaction, if the reaction is exothermic in one direction, it is endothermic in the other direction.

If the temperature is increased, the equilibrium position moves in the direction of the endothermic process. For example, sulfur dioxide reacts with oxygen in a reversible reaction:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (forward reaction is exothermic)

If the forward reaction is exothermic, the backward reaction must be endothermic. Therefore, if the temperature is increased, the equilibrium position moves to the left.