Chemistry Notes - Topics (7,8,9)

TOPIC 7: RATES OF REACTION & ENERGY CHANGES

SECTION 1: RATES OF REACTION ⏱

1.1 What is the rate of reaction?

The rate of reaction tells us how quickly reactants are used up or products are formed.

A reaction with a high rate happens quickly. A reaction with a low rate happens slowly.

Formula to learn

\boxed{\text{Rate of reaction}=\frac{\text{quantity of reactant used or product formed}}{\text{time taken}}}

Depending on the experiment, the quantity might be:

• Volume of gas produced, in cm³

• Mass lost, in g

• Amount or concentration of a reactant or product

Units depend on what is measured, for example:

• cm³/s

• g/s

• mol/dm³/s, if concentration changes are measured

Example

A reaction produces 60 cm³ of gas in 20 seconds.

\text{Rate}=\frac{60}{20}

\boxed{\text{Rate}=3\text{ cm}^3/\text{s}}

Exam tip: Always include the units in your final answer!

1.2 Collision theory 💥

Collision theory explains why reactions happen and why some reactions happen faster than others.

For a chemical reaction to occur, particles must collide with:

1. Enough energy to overcome the activation energy.

2. The correct orientation, where relevant to the reaction.

A collision that results in a reaction is called a successful collision.

Key definition: Activation energy

Activation energy is the minimum energy that reacting particles must have for a reaction to occur.

Learn this definition carefully because it can appear in a one-mark question!

Why don’t all collisions cause a reaction?

Some particles collide but do not react because they do not have enough energy to overcome the activation energy.

Even if particles collide frequently, the reaction will not be fast unless enough collisions are successful.

Collision theory summary

• More frequent collisions can increase the rate.

• A greater proportion of collisions with enough energy can increase the rate.

• More successful collisions per second means a faster reaction.

PERFECT EXAM PHRASE:

“More successful collisions occur per second, so the rate of reaction increases.”

Use this when explaining most factors that affect reaction rates—but make sure you also explain why the factor causes more successful collisions.

SECTION 2: FACTORS THAT AFFECT THE RATE OF REACTION

There are five main factors you need to know.

2.1 Concentration of a solution

Increasing the concentration of a reactant usually increases the rate of reaction.

Why?

A more concentrated solution contains more reactant particles in the same volume.

This means:

1. There are more particles close together.

2. Collisions happen more frequently.

3. There are more successful collisions per second.

4. The reaction rate increases.

Exam answer

Question: Explain why increasing the concentration of hydrochloric acid increases its rate of reaction with magnesium.

Model answer:

Increasing the concentration means there are more acid particles per unit volume. This increases the frequency of collisions between acid particles and magnesium, resulting in more successful collisions per second and a faster reaction.

2.2 Pressure of gases

Increasing the pressure of gaseous reactants can increase the rate of reaction.

Why?

When the pressure increases:

• Gas particles are compressed into a smaller volume.

• The particles are closer together.

• Collisions occur more frequently.

• There are more successful collisions per second.

• The reaction rate increases.

Important: This explanation applies to reactions involving gaseous reactants. Increasing the pressure does not automatically increase the rate of every reaction.

Exam tip

For gas-pressure questions, use the phrase “particles are closer together, so collisions occur more frequently.”

2.3 Surface area of a solid

Increasing the surface area of a solid reactant increases the rate of reaction.

For example, powdered calcium carbonate reacts faster with hydrochloric acid than the same mass of large calcium carbonate chips, assuming the other conditions are the same.

Why?

Breaking a solid into smaller pieces exposes more of its surface.

This means:

1. More solid particles are available at the surface for collisions.

2. Reactant particles can collide with the solid more frequently.

3. More successful collisions occur per second.

4. The reaction rate increases.

Remember!

Same mass + smaller pieces = larger total surface area.

The total mass hasn’t increased; the exposed surface area has.

2.4 Temperature 🌡

Increasing temperature increases the rate of reaction.

This is one of the most important collision-theory explanations.

Why?

When temperature increases:

1. Particles gain kinetic energy.

2. They move faster.

3. Collisions happen more frequently.

4. A greater proportion of particles have energy equal to or greater than the activation energy.

5. More collisions are successful per second.

6. The reaction rate increases.

Higher Tier detail ⭐

The most important effect is that a greater proportion of particles have enough energy to overcome the activation energy.

Although particles move faster and collide more frequently, the increase in the proportion of sufficiently energetic particles is particularly important.

Exam answer

Question: Explain why increasing temperature increases the rate of reaction.

Model answer:

The particles gain kinetic energy and move faster, so collisions occur more frequently. A greater proportion of particles have energy equal to or greater than the activation energy, so there are more successful collisions per second. Therefore, the rate increases.

2.5 Catalysts

A catalyst increases the rate of a chemical reaction without being used up overall.

Catalysts provide an alternative reaction pathway with a lower activation energy.

How do catalysts work?

1. The catalyst provides an alternative pathway.

2. This pathway has a lower activation energy.

3. At the same temperature, a greater proportion of particles have enough energy to react.

4. There are more successful collisions per second.

5. The reaction happens faster.

Important facts

• A catalyst is not used up overall.

• It does not change the overall energy change of the reaction.

• It lowers the activation energy.

• It can be specific to particular reactions.

Common example

Manganese dioxide can catalyse the decomposition of hydrogen peroxide, speeding up the production of oxygen.

Common exam mistake 🚨

❌ “The catalyst gives the particles more energy.”

✅ “The catalyst provides an alternative reaction pathway with a lower activation energy.”

A catalyst does not increase the particles’ kinetic energy. Temperature affects kinetic energy; a catalyst changes the reaction pathway.

2.6 Quick comparison table

Learn the explanations, not just the names of the factors. Exam questions often ask you to explain why the rate changes.

SECTION 3: MEASURING THE RATE OF REACTION 🧪

You need to understand how experiments measure reaction rates and how to interpret their results.

3.1 Measuring the volume of gas produced

A reaction that produces a gas can be monitored using a gas syringe.

For example, magnesium reacts with hydrochloric acid to produce hydrogen gas.

\boxed{\text{Mg}+2\text{HCl}\rightarrow\text{MgCl}_2+\text{H}_2}

The gas syringe measures the volume of hydrogen produced over time.

Method

1. Measure a fixed volume of hydrochloric acid into a conical flask.

2. Add a measured amount of magnesium.

3. Immediately fit the bung connected to a gas syringe.

4. Start the stopwatch.

5. Record the volume of gas at regular time intervals.

6. Repeat if necessary to improve reliability.

Safety

• Wear eye protection.

• Hydrochloric acid is an irritant at some concentrations and can be corrosive at higher concentrations.

• Hydrogen is flammable, so keep it away from naked flames.

Important practical detail: The reaction may start as soon as the magnesium is added, so work quickly and consistently to reduce measurement error.

3.2 Measuring the loss of mass

Some reactions produce a gas that escapes from the flask. As the gas leaves, the total mass measured by a balance decreases.

For example, calcium carbonate reacts with hydrochloric acid to produce carbon dioxide.

\boxed{\text{CaCO}_3+2\text{HCl}\rightarrow\text{CaCl}_2+\text{H}_2\text{O}+\text{CO}_2}

Method

1. Place a conical flask containing a measured volume of acid on a balance.

2. Add a measured amount of calcium carbonate.

3. Record the initial mass.

4. Record the mass at regular time intervals.

5. Continue until the mass stops changing significantly.

Why does the mass decrease?

Carbon dioxide gas escapes from the flask, so the measured mass of the flask and its contents decreases.

How can you calculate the mass lost?

\boxed{\text{Mass lost}=\text{initial mass}-\text{final mass}}

For example:

Initial mass = 150.0 g
Final mass = 147.6 g

\text{Mass lost}=150.0-147.6

\boxed{2.4\text{ g}}

Exam warning: If a gas is produced but cannot escape, the total mass of a closed system stays constant. The mass decreases in this experiment because gas escapes from the flask.

3.3 The disappearing-cross experiment

This experiment can investigate how concentration affects the rate of reaction between sodium thiosulfate solution and dilute hydrochloric acid.

The reaction produces solid sulfur, which makes the mixture cloudy.

As the sulfur forms, a cross underneath the flask becomes harder to see until it disappears.

Method

1. Place a conical flask on top of a black cross.

2. Add a measured volume of sodium thiosulfate solution.

3. Add dilute hydrochloric acid and start the stopwatch immediately.

4. Look down through the solution.

5. Stop the timer when the cross can no longer be seen.

6. Repeat using different concentrations of sodium thiosulfate.

Equation

A simplified overall equation is:

\boxed{\text{Na}_2\text{S}_2\text{O}_3+2\text{HCl}\rightarrow2\text{NaCl}+\text{SO}_2+\text{S}+\text{H}_2\text{O}}

Sulfur is the solid that causes the cloudiness.

Calculating an approximate rate

For this practical, a commonly used approximation is:

\boxed{\text{Relative rate}=\frac{1}{\text{time taken for the cross to disappear}}}

If the cross disappears in 40 seconds:

\text{Relative rate}=\frac{1}{40}=0.025\text{ s}^{-1}

This is a relative rate, not a direct measurement of the amount of sulfur produced per second.

Variables

• Independent variable: Concentration of sodium thiosulfate.

• Dependent variable: Time taken for the cross to disappear.

• Control variables: Temperature, total volume of solution, concentration and volume of acid, flask, cross and viewing method.

Safety

The reaction produces sulfur dioxide, which is an irritant gas. Use the required small quantities and appropriate ventilation, following your teacher’s safety instructions.

What result would you expect?

Increasing the concentration of sodium thiosulfate should make the cross disappear more quickly because there are more reactant particles per unit volume, increasing the frequency of collisions.

Evaluation point

The endpoint is judged by eye, so it is subjective. Different people may decide that the cross has disappeared at slightly different times.

To improve reliability:

• Repeat the experiment.

• Calculate a mean time.

• Keep the viewing conditions the same.

• Use the same person to judge the endpoint, if possible.

SECTION 4: RATE-OF-REACTION GRAPHS 📈

Graphs are a common source of exam marks, so make sure you understand what the shape tells you.

4.1 Volume of gas against time

Imagine a graph showing the volume of gas produced during a reaction.

• At the start, the graph rises because gas is being produced.

• The gradient is steep when the reaction is fast.

• The gradient becomes less steep as the reaction slows.

• The graph eventually becomes horizontal when no more gas is produced.

Why does the reaction slow down?

As the reaction continues, reactants are used up. Their concentration decreases, so collisions occur less frequently. The rate decreases.

Why does the graph become horizontal?

The reaction has stopped producing any more gas because the limiting reactant has been used up, or the reaction has otherwise reached its endpoint.

The final volume of gas depends on how much gas can be produced from the reactants—not simply on how quickly the reaction happens.

4.2 What does the gradient tell you?

The gradient of a graph of product formed against time tells you the rate of reaction.

For a straight section:

\boxed{\text{Gradient}=\frac{\text{change in }y}{\text{change in }x}}

For a gas-volume graph:

\boxed{\text{Rate}=\frac{\text{change in gas volume}}{\text{change in time}}}

Example

Gas volume increases from 10 cm³ to 50 cm³ between 10 seconds and 30 seconds.

\text{Rate}=\frac{50-10}{30-10}

=\frac{40}{20}

\boxed{2\text{ cm}^3/\text{s}}

Always use the change in volume divided by the change in time, not just one value from the graph.

4.3 Average rate versus rate at a particular moment

Average rate is calculated over a period of time.

\text{Average rate}=\frac{\text{change in quantity}}{\text{change in time}}

Instantaneous rate is the rate at one particular moment.

To estimate instantaneous rate from a curved graph:

1. Draw a tangent to the curve at the required point.

2. Choose two well-separated points on the tangent.

3. Calculate the gradient of the tangent.

\text{Instantaneous rate}=\text{gradient of tangent}

Exam tip ⭐

A tangent must touch the curve at the point you are investigating. Do not simply draw a line joining two points on the curve—that gives an average rate over an interval.

4.4 Comparing two rate graphs

If two reactions produce the same total amount of product:

• The steeper graph represents the faster reaction at that point.

• The graph that reaches its plateau first represents the reaction that finishes sooner.

• If the final plateaus are the same height, both reactions produced the same total amount of product.

A faster reaction does not necessarily produce more product overall.

For example, using a more concentrated acid may make a reaction faster. However, if the same amount of limiting reactant is used, the total amount of gas may stay the same.

SECTION 5: EXOTHERMIC AND ENDOTHERMIC REACTIONS 🔥❄

Now for the energy changes part!

Chemical reactions involve energy being transferred between the reacting chemicals and their surroundings.

5.1 Exothermic reactions

An exothermic reaction transfers energy to the surroundings.

The surroundings become warmer, so the temperature usually increases.

Examples

• Combustion of fuels

• Neutralisation of acids with alkalis

• Many oxidation reactions

Energy changes

During an exothermic reaction:

• More energy is released when new bonds form than is absorbed when bonds in the reactants break.

• The products have less chemical energy than the reactants.

• The energy difference is transferred to the surroundings.

For an exothermic reaction:

\boxed{\Delta H<0}

The enthalpy change is negative.

Memory trick: Exothermic = energy exits the reacting system and goes to the surroundings.

5.2 Endothermic reactions

An endothermic reaction takes in energy from the surroundings.

The surroundings become cooler, so the temperature usually decreases.

Examples

• Thermal decomposition reactions that require a continuous energy supply

• Some reactions between salts and water, such as certain dissolving processes

• Photosynthesis, which takes in energy from light

Energy changes

During an endothermic reaction:

• More energy is absorbed breaking bonds in the reactants than is released forming bonds in the products.

• The products have more chemical energy than the reactants.

• Energy is transferred from the surroundings into the reacting system.

For an endothermic reaction:

\boxed{\Delta H>0}

The enthalpy change is positive.

Memory trick: Endothermic = energy enters the reacting system from the surroundings.

5.3 Exothermic versus endothermic

Common mistake 🚨

Do not say that an exothermic reaction “creates energy.”

Energy is transferred; it is not created or destroyed.

SECTION 6: REACTION PROFILE DIAGRAMS

Reaction profile diagrams show how the energy changes as a reaction proceeds.

6.1 Exothermic reaction profile

In an exothermic reaction:

• The reactants start at a higher energy level.

• The line rises to a peak.

• The peak represents the energy needed to reach the transition state.

• The products finish at a lower energy level.

The vertical difference between the reactants and the peak represents the activation energy.

The vertical difference between the reactants and products represents the overall energy change.

6.2 Endothermic reaction profile

In an endothermic reaction:

• The reactants start at a lower energy level.

• The line rises to a peak.

• The products finish at a higher energy level.

The products have more chemical energy than the reactants because energy has been taken in overall.

6.3 What does a catalyst do to the profile?

A catalyst provides an alternative pathway with a lower activation energy.

On a reaction profile diagram, the catalysed pathway has a lower peak than the uncatalysed pathway.

However:

• The reactant energy level stays the same.

• The product energy level stays the same.

• The overall energy change stays the same.

• The catalyst lowers the activation energy.

Exam tip: A catalyst changes the pathway, not the overall energy difference between reactants and products.

SECTION 7: BOND ENERGIES — HIGHER TIER ⭐

This is a particularly important calculation area.

7.1 What happens to bonds during a reaction?

Two rules are essential:

\boxed{\text{Breaking bonds requires energy}}

\boxed{\text{Making bonds releases energy}}

Energy must be supplied to overcome the forces holding atoms together in a chemical bond.

When new bonds form, energy is released.

Remember it this way

• Break = energy in

• Make = energy out

Do not reverse these!

7.2 Calculating energy changes using bond energies

The formula is:

\boxed{\Delta H=\text{energy needed to break bonds}-\text{energy released making bonds}}

Or, in shorter form:

\boxed{\Delta H=\sum E(\text{bonds broken})-\sum E(\text{bonds made})}

Bond energies are usually given in kJ/mol.

A positive result indicates an endothermic reaction; a negative result indicates an exothermic reaction.

7.3 Worked example: exothermic reaction

Suppose:

• Energy needed to break the reactant bonds = 500 kJ/mol

• Energy released when product bonds form = 700 kJ/mol

Calculate the enthalpy change.

Step 1: Write the formula

\Delta H=\text{energy to break bonds}-\text{energy released making bonds}

Step 2: Substitute

\Delta H=500-700

Step 3: Calculate

\boxed{\Delta H=-200\text{ kJ/mol}}

The answer is negative, so the reaction is exothermic.

Why? More energy is released forming bonds than is absorbed breaking bonds.

7.4 Worked example: endothermic reaction

Suppose:

• Energy needed to break the reactant bonds = 850 kJ/mol

• Energy released when product bonds form = 600 kJ/mol

\Delta H=850-600

\boxed{\Delta H=+250\text{ kJ/mol}}

The answer is positive, so the reaction is endothermic.

Why? More energy is absorbed breaking bonds than is released forming bonds.

7.5 Bond-energy calculation method

Use this four-step method every time:

Step 1: Identify all the bonds broken in the reactants.

Step 2: Identify all the bonds made in the products.

Step 3: Add the energies for each set of bonds, including the number of each type of bond.

Step 4: Calculate:

\Delta H=\text{total energy of bonds broken}-\text{total energy of bonds made}

Common exam mistake

If the equation contains two molecules with the same bond, remember to count all of those bonds.

For example, two molecules of hydrogen contain two H–H bonds in total, not one.

Also, use the correct bond energy from the question’s data table. Bond-energy values are often average values.

SECTION 8: CALORIMETRY AND TEMPERATURE CHANGES 🧪

Experiments can estimate energy changes by measuring temperature changes.

For example, a student might mix an acid with an alkali and record the temperature before and after the reaction.

Basic method

1. Measure a known volume of solution into an insulated container.

2. Measure and record the initial temperature.

3. Add the second reactant.

4. Stir carefully.

5. Record the highest or lowest temperature reached, as appropriate.

6. Calculate the temperature change.

\boxed{\text{Temperature change}=\text{final temperature}-\text{initial temperature}}

Example

Initial temperature = 19°C
Final temperature = 26°C

\Delta T=26-19

\boxed{\Delta T=+7^\circ\text{C}}

The temperature increased, suggesting the reaction was exothermic.

If the temperature decreased, this would suggest an endothermic process.

Improving the experiment

Heat can be transferred to or from the surroundings, so the measured temperature change may be smaller than the true temperature change.

Improvements include:

• Use an insulated container.

• Use a lid to reduce heat loss.

• Stir to make the temperature more uniform.

• Use a temperature probe or thermometer carefully.

• Keep the volumes and concentrations controlled.

Important: A temperature increase suggests energy transfer to the surroundings; a temperature decrease suggests energy transfer from the surroundings. Interpret the result in the context of the experiment.

TOPIC 8: FUELS, EARTH AND THE ATMOSPHERE

SECTION 1: CRUDE OIL 🛢

1.1 What is crude oil?

Crude oil is a finite resource made from the remains of ancient organisms.

It is:

• A mixture of many different hydrocarbons.

• A non-renewable resource.

• Used as a source of fuels and as a feedstock for making chemicals.

What is a hydrocarbon?

A hydrocarbon is a compound containing hydrogen and carbon atoms only.

Examples:

• Methane: CH₄

• Ethane: C₂H₆

• Propane: C₃H₈

• Butane: C₄H₁₀

Crude oil contains hydrocarbons with different chain lengths and properties.

How was crude oil formed?

Millions of years ago:

1. Tiny marine organisms died and settled on the seabed.

2. Their remains became buried beneath layers of sediment.

3. Over millions of years, heat and pressure transformed this organic material into crude oil and natural gas.

Exam tip: Remember that crude oil formed over millions of years, so it cannot be replaced as quickly as we use it.

1.2 Why is crude oil a mixture?

Crude oil contains many different hydrocarbons mixed together.

These hydrocarbons have different:

• Boiling points

• Chain lengths

• Viscosities

• Flammabilities

• Volatilities

Because the hydrocarbons have different boiling points, they can be separated by fractional distillation.

SECTION 2: FRACTIONAL DISTILLATION 🔥

2.1 What is fractional distillation?

Fractional distillation separates a mixture of liquids with different boiling points.

It is used to separate crude oil into useful fractions.

A fraction is a mixture of hydrocarbons with similar boiling points and similar numbers of carbon atoms.

2.2 How does fractional distillation work?

Learn these steps in order:

Step 1: Crude oil is heated until most of it vaporises.

Step 2: The vapours enter a fractionating column.

Step 3: The column is hottest at the bottom and coolest at the top.

Step 4: The vapours rise and cool.

Step 5: Different hydrocarbons condense at different temperatures.

Step 6: Fractions are collected at different heights in the column.

Why do the fractions separate?

Different hydrocarbons have different boiling points.

• High-boiling-point hydrocarbons condense lower down, where it is hotter.

• Low-boiling-point hydrocarbons rise further and condense higher up, where it is cooler.

• The smallest, lowest-boiling-point hydrocarbons may leave at the top as gases.

Important: Fractional distillation is a physical separation process. It does not chemically change the hydrocarbons.

2.3 Trends in the fractions

As hydrocarbon chain length increases:

What do these words mean?

Viscosity: How resistant a liquid is to flowing.

• Low viscosity: flows easily.

• High viscosity: flows slowly and is thick.

Volatility: How easily a substance evaporates.

• High volatility: evaporates easily.

• Low volatility: does not evaporate easily.

Why do longer hydrocarbons have higher boiling points?

Longer hydrocarbon molecules have more electrons and larger electron clouds, so the intermolecular forces between their molecules are stronger.

More energy is needed to overcome these forces, giving them higher boiling points.

Higher Tier note: These are intermolecular forces between molecules—not the covalent bonds within the hydrocarbon molecules. Boiling does not normally break the covalent bonds within the molecules.

2.4 What are the fractions used for?

You do not always need to memorise every fraction in this exact order, but you should understand that different fractions have different uses because their properties differ.

Exam-style question

Why is petrol suitable for use as a car fuel?

A good answer could mention that petrol contains relatively short-chain hydrocarbons that are volatile and ignite relatively easily, making them useful in engines.

SECTION 3: ALKANES AND THEIR FORMULAE 🧪

3.1 What are alkanes?

Alkanes are saturated hydrocarbons.

They contain:

• Carbon and hydrogen only.

• Single covalent bonds between carbon atoms.

• The maximum possible number of hydrogen atoms for their carbon skeleton.

The general formula for alkanes is:

\boxed{C_nH_{2n+2}}

Here, n represents the number of carbon atoms.

The first four alkanes

How to work out an alkane formula

For an alkane with 5 carbon atoms:

n=5

Using the general formula:

C_nH_{2n+2}

C_5H_{(2\times5)+2}

\boxed{C_5H_{12}}

This is pentane.

Quick practice

What is the formula of an alkane with 7 carbon atoms?

C_nH_{2n+2}

C_7H_{16}

Answer: C₇H₁₆.

SECTION 4: COMBUSTION OF HYDROCARBONS 🔥

Hydrocarbons are commonly burned as fuels because combustion releases energy.

4.1 Complete combustion

Complete combustion occurs when a hydrocarbon burns in a plentiful supply of oxygen.

The products are:

• Carbon dioxide

• Water

General word equation

Hydrocarbon + oxygen → carbon dioxide + water

Example: methane

\boxed{CH_4+2O_2\rightarrow CO_2+2H_2O}

This equation is balanced:

• Carbon: 1 on each side

• Hydrogen: 4 on each side

• Oxygen: 4 on each side

Complete combustion releases energy that can be used for heating, transport or electricity generation.

4.2 Incomplete combustion

Incomplete combustion occurs when there is insufficient oxygen for complete combustion.

It can produce:

• Carbon monoxide

• Carbon (soot)

• Water

The products depend on how much oxygen is available.

Example: formation of carbon monoxide

\boxed{2CH_4+3O_2\rightarrow2CO+4H_2O}

Example: formation of carbon

\boxed{CH_4+O_2\rightarrow C+2H_2O}

These are balanced equations for possible incomplete-combustion reactions.

Why is carbon monoxide dangerous?

Carbon monoxide is:

• Colourless

• Odourless

• Toxic

It binds to haemoglobin in red blood cells, reducing the blood’s ability to carry oxygen around the body.

This can cause serious harm or death.

Exam tip: If asked why incomplete combustion is dangerous, mention toxic carbon monoxide, rather than simply saying it produces less energy.

Why does soot cause problems?

Soot consists of tiny carbon particles. These can contribute to air pollution and can be harmful when inhaled.

SECTION 5: CRACKING HYDROCARBONS ⚗

5.1 Why is cracking needed?

Crude oil contains a mixture of hydrocarbons, but the demand for different fractions is not equal.

For example:

• Shorter-chain hydrocarbons are useful as fuels.

• Some longer-chain hydrocarbons are less useful and may be in excess.

• The chemical industry needs alkenes to manufacture polymers and other chemicals.

Cracking breaks long-chain hydrocarbons into smaller, more useful molecules.

It can produce:

• Shorter-chain alkanes

• Alkenes

5.2 What is an alkene?

Alkenes are unsaturated hydrocarbons containing at least one carbon–carbon double bond.

The general formula for alkenes with one double bond and one carbon–carbon chain is:

\boxed{C_nH_{2n}}

Examples:

• Ethene: C₂H₄

• Propene: C₃H₆

• Butene: C₄H₈

The double bond makes alkenes more reactive than alkanes in many common reactions.

5.3 How does cracking work?

A typical cracking process involves:

1. Vaporising the long-chain hydrocarbon.

2. Passing the vapour over a hot catalyst, or heating it strongly under suitable conditions.

3. Breaking some carbon–carbon bonds.

4. Producing smaller hydrocarbons.

The precise conditions depend on the cracking process.

Example equation

\boxed{C_{10}H_{22}\rightarrow C_8H_{18}+C_2H_4}

Decane breaks down to form octane and ethene.

Check the atoms:

• Carbon: 10=8+2

• Hydrogen: 22=18+4

The equation is balanced.

Another possible example

\boxed{C_{12}H_{26}\rightarrow C_8H_{18}+C_4H_8}

Dodecane breaks down into octane and butene.

Why is cracking useful?

• It converts less useful long-chain hydrocarbons into more useful shorter-chain hydrocarbons.

• It produces alkenes, which are used to manufacture polymers.

• It helps meet demand for particular fuels and chemical feedstocks.

Common exam mistake 🚨

Do not say cracking is the same as fractional distillation.

SECTION 6: THE EARLY ATMOSPHERE 🌋

Now we’re moving from fuels to the history of Earth!

6.1 What was Earth’s early atmosphere like?

Scientists use evidence to develop theories about Earth’s early atmosphere.

The early atmosphere is thought to have contained large amounts of:

• Carbon dioxide

• Water vapour

It may also have contained:

• Nitrogen

• Small amounts of other gases, including methane and ammonia

The exact composition of the earliest atmosphere is uncertain.

Important: Earth’s atmosphere has changed substantially over billions of years.

6.2 How did Earth’s early atmosphere form?

One scientific explanation is that gases were released by volcanic activity.

This process is called volcanic outgassing.

As Earth cooled:

1. Water vapour in the atmosphere condensed.

2. Liquid water formed.

3. Water collected to form the oceans.

4. Carbon dioxide dissolved in the oceans and became involved in chemical processes that removed it from the atmosphere.

The atmosphere continued to change through geological, chemical and biological processes.

SECTION 7: HOW OXYGEN INCREASED 🌱

7.1 The development of photosynthetic organisms

Early photosynthetic organisms, including ancient bacteria, began releasing oxygen through photosynthesis.

The word equation for photosynthesis is:

\boxed{\text{Carbon dioxide}+\text{water}\rightarrow\text{glucose}+\text{oxygen}}

Light energy is required.

As photosynthetic organisms became more widespread, oxygen gradually accumulated in the atmosphere.

Why did oxygen not build up immediately?

Initially, oxygen reacted with other substances, including dissolved iron and rocks.

It took a very long time for oxygen to accumulate significantly in the atmosphere.

Key exam idea

Photosynthesis was a major cause of the increase in atmospheric oxygen over geological time.

SECTION 8: HOW CARBON DIOXIDE DECREASED 🌊

Carbon dioxide levels decreased through several processes.

8.1 Dissolving in the oceans

Carbon dioxide dissolves in water.

Some dissolved carbon dioxide then participates in chemical reactions, including reactions that form carbonate compounds.

8.2 Formation of carbonate rocks

Marine organisms can use dissolved substances to make shells and skeletons containing calcium carbonate.

When these organisms die, their remains can become buried and eventually contribute to sedimentary rocks such as limestone.

This stores carbon for long periods.

8.3 Formation of fossil fuels

Some ancient organisms were buried and, under suitable conditions over millions of years, contributed to the formation of fossil fuels.

When these fuels are burned, carbon that was stored underground is released as carbon dioxide.

8.4 Photosynthesis

Photosynthesis removes carbon dioxide from the atmosphere.

Plants and algae use it to make glucose.

Summary: why did atmospheric carbon dioxide decrease?

• It dissolved in oceans.

• Carbon became locked up in carbonate rocks.

• Carbon became stored in fossil fuels.

• Photosynthetic organisms removed it from the atmosphere.

SECTION 9: THE MODERN ATMOSPHERE 🌍

Today’s atmosphere consists mainly of nitrogen and oxygen.

Approximate composition of dry air:

These are approximate values; atmospheric carbon dioxide changes over time.

What is nitrogen used for?

Nitrogen is relatively unreactive under ordinary conditions and makes up most of the atmosphere.

What is oxygen used for?

Oxygen is needed for aerobic respiration and supports combustion.

Why is carbon dioxide important?

Carbon dioxide is:

• Used by plants in photosynthesis.

• A greenhouse gas.

• Released by respiration, combustion and some industrial processes.

Exam tip: Do not confuse oxygen with ozone. Oxygen is O₂; ozone is O₃.

SECTION 10: GREENHOUSE GASES AND THE GREENHOUSE EFFECT ☀🌡

This is one of the most important parts of the topic.

10.1 What is the greenhouse effect?

The greenhouse effect is a natural process that keeps Earth’s surface warmer than it would otherwise be.

How does it work?

1. Shortwave radiation from the Sun passes through the atmosphere.

2. Earth’s surface absorbs some of this energy and warms up.

3. The warm surface emits infrared radiation.

4. Greenhouse gases absorb some of this outgoing infrared radiation.

5. These gases emit infrared radiation in different directions, including back towards Earth’s surface.

6. This reduces the rate at which energy escapes to space, warming the surface and lower atmosphere.

Greenhouse gases include:

• Carbon dioxide

• Methane

• Water vapour

• Nitrous oxide

Important distinction

The greenhouse effect is natural and necessary for life as we know it.

However, increasing greenhouse gas concentrations strengthens the greenhouse effect and contributes to global warming.

10.2 What is the enhanced greenhouse effect?

Human activities have increased the atmospheric concentrations of several greenhouse gases.

Examples include:

• Burning fossil fuels, which releases carbon dioxide.

• Deforestation, which reduces carbon dioxide removal by photosynthesis and can release stored carbon.

• Agriculture, which releases methane and nitrous oxide.

• Some industrial processes, which release greenhouse gases.

These changes enhance the greenhouse effect, causing additional warming.

Exam-style explanation

Question: Explain how burning fossil fuels contributes to global warming.

Model answer:

Burning fossil fuels releases carbon dioxide into the atmosphere. Carbon dioxide absorbs some of the infrared radiation emitted by Earth’s surface and emits infrared radiation in different directions. Increasing its concentration strengthens the greenhouse effect, increasing Earth’s average temperature.

For a stronger answer, explain the energy transfer, not just that carbon dioxide is a greenhouse gas.

SECTION 11: GLOBAL WARMING AND CLIMATE CHANGE 🌡

11.1 What is global warming?

Global warming is the long-term increase in Earth’s average surface temperature.

It is primarily driven by the enhanced greenhouse effect caused by human activities.

11.2 What is climate change?

Climate change refers to long-term changes in average weather conditions and patterns.

It can include changes in:

• Temperature

• Rainfall

• Frequency or intensity of some extreme weather events

• Sea levels

• Ice cover

Global warming contributes to climate change.

11.3 Possible effects of climate change

Examples include:

• Melting glaciers and ice sheets.

• Rising sea levels due to melting land ice and the expansion of warmer seawater.

• Changes in rainfall patterns.

• Increased risks of heatwaves in many regions.

• Changes in habitats and the distribution of species.

• Impacts on food production and water availability.

The effects vary by location, and the severity depends partly on how much further the climate changes.

Exam tip

When asked for an effect of climate change, give a clear cause-and-effect link.

For example:

“Rising temperatures can melt land ice, adding water to the oceans and contributing to sea-level rise.”

That is stronger than simply writing “the ice melts.”

SECTION 12: ATMOSPHERIC POLLUTION 🚗

Burning fuels can release pollutants, especially when combustion is incomplete or when fuels contain impurities.

12.1 Carbon monoxide

Source: Incomplete combustion of carbon-containing fuels.

Why it is harmful: It binds to haemoglobin and reduces the blood’s ability to carry oxygen.

Prevention: Ensure appliances are properly installed and maintained, and that combustion systems have suitable air supplies.

12.2 Sulfur dioxide

Source: Burning fuels containing sulfur impurities.

Why it is harmful:

• It irritates the respiratory system.

• It contributes to acid rain.

How does acid rain form?

Sulfur dioxide can react in the atmosphere to form acids. These acidic substances can dissolve in rainwater and fall to Earth.

Acid rain can:

• Damage plants and aquatic ecosystems.

• Acidify lakes and soils.

• Weather limestone buildings and statues.

Some nitrogen oxides also contribute to acid rain.

12.3 Nitrogen oxides

At high temperatures inside engines, nitrogen and oxygen from the air can react to form nitrogen oxides.

These gases can:

• Irritate the lungs.

• Contribute to acid rain.

• Contribute to the formation of photochemical smog.

Why do engines produce nitrogen oxides?

At high engine temperatures, nitrogen and oxygen can react even though they are relatively unreactive at ordinary temperatures.

12.4 Particulates

Particulates are tiny solid particles or liquid droplets suspended in the air.

They can come from incomplete combustion and other sources.

Why are they harmful?

Small particles can enter the respiratory system and worsen breathing problems.

Pollution summary table

Exam tip

Don’t mix up the pollutants!

• Carbon monoxide: poisonous because it affects haemoglobin.

• Carbon dioxide: greenhouse gas.

• Sulfur dioxide: contributes to acid rain.

• Nitrogen oxides: contribute to acid rain and smog.

• Particulates: damage health, especially the respiratory system.

SECTION 13: REDUCING POLLUTION AND CARBON EMISSIONS 🌱

There are several ways to reduce emissions from energy use.

13.1 Renewable energy

Examples include:

• Wind

• Solar

• Hydroelectric power

• Tidal power

These energy sources do not require burning fossil fuels to generate electricity during normal operation.

However, building and maintaining renewable-energy systems can still produce emissions.

13.2 Energy efficiency

Using less energy for the same purpose reduces fuel demand.

Examples:

• Insulating buildings.

• Using efficient appliances.

• Improving vehicle efficiency.

• Reducing unnecessary energy use.

13.3 Public transport and electric vehicles

Public transport can reduce emissions per passenger compared with many individual car journeys.

Electric vehicles have no exhaust emissions while driving, but their overall emissions depend partly on how the electricity is generated and how the vehicles and batteries are manufactured.

13.4 Carbon capture and storage

Carbon capture and storage aims to capture carbon dioxide from some industrial processes or power generation and store it rather than releasing it into the atmosphere.

It may reduce emissions from suitable sources, but it does not eliminate all greenhouse gas emissions.

13.5 Planting and protecting forests

Plants absorb carbon dioxide during photosynthesis.

Protecting forests and restoring suitable ecosystems can help remove carbon dioxide from the atmosphere and store carbon in vegetation and soils.

Exam tip: If asked how a country could reduce its carbon emissions, explain how your chosen method reduces the amount of carbon dioxide released or increases its removal.

SECTION 14: KEY EQUATIONS TO MEMORISE 🧠

These are useful equations for this topic.

Complete combustion of methane

\boxed{CH_4+2O_2\rightarrow CO_2+2H_2O}

Incomplete combustion producing carbon monoxide

\boxed{2CH_4+3O_2\rightarrow2CO+4H_2O}

Incomplete combustion producing carbon

\boxed{CH_4+O_2\rightarrow C+2H_2O}

Example of cracking

\boxed{C_{10}H_{22}\rightarrow C_8H_{18}+C_2H_4}

Photosynthesis word equation

\boxed{\text{Carbon dioxide}+\text{water}\rightarrow\text{glucose}+\text{oxygen}}

Remember that light energy is required for photosynthesis.

TOPIC 9: QUALITATIVE ANALYSIS, ORGANIC CHEMISTRY & NANOPARTICLES

SECTION 1: QUALITATIVE ANALYSIS 🧪

1.1 What is qualitative analysis?

Qualitative analysis is used to identify which substances or ions are present in a sample.

For example, you might be given an unknown solution and asked to identify whether it contains:

• Copper(II) ions

• Iron(II) ions

• Iron(III) ions

• Carbonate ions

• Sulfate ions

• Halide ions

You identify substances by carrying out chemical tests and observing the results.

Key terms

• Observation: What you see, such as a colour change or precipitate.

• Inference: What you conclude from the observation.

• Precipitate: An insoluble solid formed when solutions react.

Exam tip: State the observation first, then explain what it tells you.

For example: “A blue precipitate forms, indicating the presence of copper(II) ions.”

SECTION 2: TESTS FOR POSITIVE IONS — CATIONS

A cation is a positively charged ion.

You need to know how aqueous metal ions react with sodium hydroxide solution.

2.1 Sodium hydroxide tests

Adding sodium hydroxide solution to a sample containing certain metal ions can produce a coloured precipitate.

Learn this table carefully!l;

Important: A white precipitate alone does not distinguish aluminium ions from calcium ions. Additional tests or evidence may be needed.

Memory trick 💗

• Copper = blue

• Iron(II) = green

• Iron(III) = brown

• Aluminium and calcium = white

2.2 What is a precipitate?

A precipitate is an insoluble solid that forms when two solutions react.

For example, when sodium hydroxide is added to a solution containing copper(II) ions, a blue precipitate of copper(II) hydroxide forms.

The ionic equation is:

\boxed{\mathrm{Cu^{2+}(aq)+2OH^-(aq)\rightarrow Cu(OH)_2(s)}}

The symbol (s) means solid.

The symbol (aq) means dissolved in water.

Another example: iron(II)

\boxed{\mathrm{Fe^{2+}(aq)+2OH^-(aq)\rightarrow Fe(OH)_2(s)}}

This produces a green precipitate.

Iron(III)

\boxed{\mathrm{Fe^{3+}(aq)+3OH^-(aq)\rightarrow Fe(OH)_3(s)}}

This produces a brown precipitate.

Higher Tier tip: If your exam asks for an ionic equation, include the correct charges and balance both atoms and electrical charge.

SECTION 3: TESTS FOR NEGATIVE IONS — ANIONS

An anion is a negatively charged ion.

You need to know tests for carbonate, sulfate and halide ions.

3.1 Carbonate ions

The carbonate ion is:

\boxed{\mathrm{CO_3^{2-}}}

Test

1. Add a dilute acid to the sample.

2. If carbonate ions are present, effervescence occurs because carbon dioxide gas is produced.

3. Bubble the gas through limewater.

4. Limewater turns milky or cloudy if carbon dioxide is present.

Ionic equation

\boxed{\mathrm{CO_3^{2-}+2H^+\rightarrow CO_2+H_2O}}

Gas test

Carbon dioxide turns limewater milky.

The cloudiness is caused by the formation of a calcium carbonate precipitate.

Exam answer

Question: Describe how you would test for carbonate ions.

Model answer: Add dilute acid to the sample. If carbonate ions are present, carbon dioxide is released. Bubble the gas through limewater; the limewater turns milky.

3.2 Sulfate ions

The sulfate ion is:

\boxed{\mathrm{SO_4^{2-}}}

Test

1. Acidify the sample with dilute hydrochloric acid.

2. Add barium chloride solution.

3. A white precipitate forms if sulfate ions are present.

The precipitate is barium sulfate.

Ionic equation

\boxed{\mathrm{Ba^{2+}(aq)+SO_4^{2-}(aq)\rightarrow BaSO_4(s)}}

Why acidify first?

Acidifying the sample helps remove interference from certain ions, particularly carbonate ions, which could otherwise produce a precipitate and lead to an incorrect conclusion.

Remember: Sulfate test = acidify, then add barium chloride = white precipitate.

3.3 Halide ions

Halides are ions formed by halogens.

The main ones you need to recognise are:

Test

1. Acidify the sample with dilute nitric acid.

2. Add silver nitrate solution.

3. Observe the precipitate colour.

Ionic equations

Chloride:

\boxed{\mathrm{Ag^+(aq)+Cl^-(aq)\rightarrow AgCl(s)}}

Bromide:

\boxed{\mathrm{Ag^+(aq)+Br^-(aq)\rightarrow AgBr(s)}}

Iodide:

\boxed{\mathrm{Ag^+(aq)+I^-(aq)\rightarrow AgI(s)}}

Why use nitric acid?

Nitric acid does not introduce chloride ions, unlike hydrochloric acid.

If you used hydrochloric acid, the chloride ions could react with silver ions and produce a white precipitate, interfering with the test.

Memory trick

• Chloride = white

• Bromide = cream

• Iodide = yellow

SECTION 4: TESTS FOR GASES 💨

These are extremely important for practical questions!

4.1 Hydrogen

Hydrogen is identified using a lit splint.

If hydrogen is present, it burns with a squeaky pop.

Equation:

\boxed{2H_2+O_2\rightarrow2H_2O}

This is the combustion of hydrogen.

4.2 Oxygen

Oxygen supports combustion.

A glowing splint relights when placed in a sample containing sufficient oxygen.

Do not confuse this with hydrogen: hydrogen gives a squeaky pop with a lit splint.

4.3 Carbon dioxide

Bubble the gas through limewater.

A positive result is that the limewater turns milky.

The reaction can be represented as:

\boxed{\mathrm{Ca(OH)_2+CO_2\rightarrow CaCO_3+H_2O}}

The calcium carbonate forms a white precipitate.

4.4 Chlorine

Chlorine is a toxic gas, so it must only be tested using appropriate small-scale laboratory procedures and suitable safety precautions.

Damp blue litmus paper turns red and is then bleached white.

Exam tip: If a question asks for the test and result, you need both. Writing “use limewater” without stating that it turns milky may not earn all the marks.

SECTION 5: ORGANIC CHEMISTRY — HYDROCARBONS 🛢

Organic chemistry is the study of carbon-containing compounds, particularly hydrocarbons and their derivatives.

5.1 Alkanes

Alkanes are saturated hydrocarbons containing only single carbon–carbon bonds.

Their general formula is:

\boxed{C_nH_{2n+2}}

Examples:

Why are they called saturated?

They contain the maximum possible number of hydrogen atoms for their carbon skeleton.

5.2 Alkenes

Alkenes are unsaturated hydrocarbons containing a carbon–carbon double bond.

For alkenes with one double bond and one carbon–carbon chain, the general formula is:

\boxed{C_nH_{2n}}

Examples:

What does unsaturated mean?

It means the molecule contains a carbon–carbon double bond and can undergo addition reactions across that bond.

5.3 The bromine water test

Bromine water can distinguish between an alkane and an alkene under suitable conditions.

Method

Add bromine water to the hydrocarbon sample and shake gently.

Why?

Alkenes react with bromine across their carbon–carbon double bond. Alkanes do not normally react with bromine water under these conditions.

Example equation

Ethene reacts with bromine to form dibromoethane:

\boxed{C_2H_4+Br_2\rightarrow C_2H_4Br_2}

This is an addition reaction because the bromine atoms add across the double bond.

Exam tip: Orange to colourless = evidence of an alkene.

SECTION 6: POLYMERS 🧵

6.1 What is a polymer?

A polymer is a large molecule made from many repeating smaller units called monomers.

A monomer is a small molecule that can join with other monomers to form a polymer.

Examples:

• Ethene is a monomer.

• Poly(ethene) is a polymer.

Polymers are used to make products such as packaging, pipes, clothing fibres and containers.

6.2 Addition polymerisation

Addition polymerisation occurs when many unsaturated monomers join together to form a polymer.

The monomers contain carbon–carbon double bonds.

During the reaction:

1. The double bond in each monomer opens.

2. The monomers join together.

3. A long polymer chain forms.

4. No small molecule is eliminated as a by-product.

Example: ethene to poly(ethene)

\boxed{n\,CH_2{=}CH_2\rightarrow[-CH_2-CH_2-]_n}

The n means that many ethene molecules join together.

Important facts

• Monomer: ethene.

• Polymer: poly(ethene).

• Type of reaction: addition polymerisation.

• The carbon–carbon double bond becomes a single bond in the polymer backbone.

Exam tip

When drawing the repeating unit of poly(ethene), show the bonds extending from both ends of the repeating unit to indicate that the chain continues.

6.3 Environmental problems caused by plastics

Many plastics are useful because they are lightweight, durable and inexpensive. However, plastic waste can cause environmental problems.

Examples:

• Many plastics take a long time to break down.

• Plastic waste can accumulate in the environment.

• Animals may become entangled in plastic or mistake it for food.

• Tiny plastic particles can contaminate ecosystems.

• Burning some plastics can release harmful substances.

Ways to reduce the problems

• Reuse plastic products where appropriate.

• Recycle suitable plastics.

• Reduce unnecessary single-use plastics.

• Develop materials that are easier to reuse or manage safely.

Exam tip: Recycling can conserve raw materials and reduce waste, but recycling systems require energy and do not necessarily eliminate all environmental impacts.

SECTION 7: ALCOHOLS 🍶

Alcohols are organic compounds containing the –OH functional group.

A functional group is the atom or group of atoms responsible for a characteristic pattern of reactions.

7.1 The first four alcohols

For these straight-chain alcohols with one –OH group, the molecular formula changes as the carbon chain gets longer.

7.2 Uses of alcohols

Alcohols have several uses, including:

• Solvents.

• Fuels.

• Chemical feedstocks.

• Ethanol in alcoholic drinks for adults.

For GCSE Chemistry, the important point is that alcohols have a range of physical and chemical properties because of their functional group and carbon chain.

7.3 Combustion of ethanol

Ethanol burns in oxygen to produce carbon dioxide and water under complete-combustion conditions.

\boxed{C_2H_5OH+3O_2\rightarrow2CO_2+3H_2O}

Combustion releases energy, which is why ethanol can be used as a fuel.

SECTION 8: CARBOXYLIC ACIDS 🧪

Carboxylic acids contain the –COOH functional group.

The first few straight-chain carboxylic acids include:

Ethanoic acid is the acid found in vinegar, although vinegar is a solution containing ethanoic acid in water.

8.1 Properties of carboxylic acids

Carboxylic acids:

• Are weak acids in aqueous solution.

• Turn blue litmus paper red.

• React with some metals to produce a salt and hydrogen.

• React with carbonates to produce a salt, water and carbon dioxide.

• React with alcohols to form esters under suitable conditions.

What does weak acid mean?

A weak acid only partially ionises in water.

This differs from a strong acid, which ionises completely in aqueous solution.

Important: Weak does not simply mean dilute. Strength describes the extent of ionisation, whereas concentration describes how much solute is present in a given volume.

8.2 Carboxylic acids reacting with metals

For example, ethanoic acid reacts with magnesium to form magnesium ethanoate and hydrogen.

A word equation is:

Ethanoic acid + magnesium → magnesium ethanoate + hydrogen

The hydrogen can be tested with a lit splint and gives a squeaky pop.

8.3 Carboxylic acids reacting with carbonates

A carboxylic acid reacts with a carbonate to produce:

• A salt

• Water

• Carbon dioxide

For example:

Ethanoic acid + sodium carbonate → sodium ethanoate + water + carbon dioxide

The carbon dioxide turns limewater milky.

SECTION 9: ESTERS 🌸

9.1 What is an ester?

Esters are organic compounds that can be formed by reacting an alcohol with a carboxylic acid.

They often have fruity or sweet smells and are used in:

• Perfumes.

• Food flavourings.

• Solvents.

9.2 How are esters made?

An alcohol reacts with a carboxylic acid, usually with an acid catalyst such as concentrated sulfuric acid.

The reaction is called esterification.

\boxed{\text{Alcohol}+\text{carboxylic acid}\rightleftharpoons\text{ester}+\text{water}}

Example: ethanol and ethanoic acid

Ethanol reacts with ethanoic acid to form ethyl ethanoate and water.

\boxed{C_2H_5OH+CH_3COOH\rightleftharpoons CH_3COOC_2H_5+H_2O}

The ester is ethyl ethanoate.

How do ester names work?

The name has two parts:

1. The first part comes from the alcohol.

2. The second part comes from the carboxylic acid.

For example:

• Ethanol gives the ethyl part.

• Ethanoic acid gives the ethanoate part.

Therefore, the ester is ethyl ethanoate.

Safety note

In a school practical, esterification may involve heating and concentrated sulfuric acid, which is corrosive. Follow your teacher’s instructions and use appropriate protective equipment.

SECTION 10: BULK AND SURFACE PROPERTIES

This section explores why materials have different properties and how those properties affect their uses.

10.1 What are bulk properties?

Bulk properties are properties of a material that are measured for the material as a whole.

Examples include:

• Density.

• Melting point.

• Electrical conductivity.

• Thermal conductivity.

• Strength.

A material’s structure and bonding influence its properties.

For example, metals generally conduct electricity because they contain delocalised electrons that can move through the structure.

10.2 What are surface properties?

Surface properties describe how a material behaves at its surface.

Examples can include:

• How readily it reacts with other substances.

• How well it adsorbs substances.

• How it interacts with liquids or gases.

Adsorption is when particles collect on the surface of a material.

Do not confuse adsorption with absorption:

• Adsorption: particles collect on a surface.

• Absorption: a substance enters the bulk of another material.

SECTION 11: NANOPARTICLES 🔬

11.1 What are nanoparticles?

Nanoparticles are extremely small particles, typically with dimensions between approximately 1 and 100 nanometres.

One nanometre is:

\boxed{1\text{ nm}=10^{-9}\text{ m}}

A nanometre is one billionth of a metre.

Nanoparticles contain relatively few atoms compared with larger particles of the same substance.

11.2 Why do nanoparticles have different properties?

Nanoparticles have a very high surface-area-to-volume ratio compared with larger particles of the same material.

This means a larger proportion of their atoms are at or near the surface.

As a result, nanoparticles may have different properties from the same substance in larger pieces.

Depending on the material, these differences can affect:

• Chemical reactivity.

• Catalytic activity.

• Optical properties.

• Electrical properties.

Example

A catalyst made from nanoparticles may have a greater exposed surface area than the same mass of the material in larger particles.

This can provide more sites where reactions occur and increase the catalyst’s effectiveness.

11.3 Uses of nanoparticles

Nanoparticles have a variety of uses.

The precise benefit depends on the material and how the nanoparticles are used.

11.4 Risks of nanoparticles

Nanoparticles may pose risks because:

• Their small size can allow some particles to enter the body through inhalation or other routes.

• Some may interact with cells differently from larger particles of the same substance.

• Their long-term environmental and health effects are not fully understood for every type of nanoparticle.

Should nanoparticles be used?

There are potential benefits and potential risks.

Scientists need to assess:

• Toxicity.

• Exposure routes.

• Environmental persistence.

• Effects on living organisms.

• Whether the benefits outweigh the risks.

Exam tip: Avoid claiming that all nanoparticles are dangerous or that all are safe. Their effects depend on the material, size, shape, coating and exposure.

SECTION 12: KEY EQUATIONS TO MEMORISE 🧠

Here are the most useful equations from this part.

Carbonate ions reacting with acid

\boxed{\mathrm{CO_3^{2-}+2H^+\rightarrow CO_2+H_2O}}

Sulfate-ion test

\boxed{\mathrm{Ba^{2+}+SO_4^{2-}\rightarrow BaSO_4(s)}}

Chloride-ion test

\boxed{\mathrm{Ag^++Cl^-\rightarrow AgCl(s)}}

Bromide-ion test

\boxed{\mathrm{Ag^++Br^-\rightarrow AgBr(s)}}

Iodide-ion test

\boxed{\mathrm{Ag^++I^-\rightarrow AgI(s)}}

Combustion of hydrogen

\boxed{2H_2+O_2\rightarrow2H_2O}

Combustion of ethanol

\boxed{C_2H_5OH+3O_2\rightarrow2CO_2+3H_2O}

Addition of bromine to ethene

\boxed{C_2H_4+Br_2\rightarrow C_2H_4Br_2}

Esterification

\boxed{\text{Alcohol}+\text{carboxylic acid}\rightleftharpoons\text{ester}+\text{water}}