teaching-notes — Chemistry (CHEMICAL REACTIONS)

Topic: CHEMICAL REACTIONS

Subtopic: Rates of chemical reactions

AS SPEED OF A CHEMICAL REACTION

In chemistry, different reactions take place at different speeds. Some reactions, like the rusting of iron, are very slow and take weeks or months. Other reactions, like the explosion of fireworks or the neutralization of an acid by an alkali, happen in a fraction of a second. The speed at which a chemical reaction takes place is called the rate of reaction.

Definition of Rate of Reaction

The rate of a chemical reaction is defined as the change in the concentration or amount of reactants or products per unit time.

To find the rate of a chemical reaction, we can measure either:

• How fast a reactant is used up (decreases in mass or volume).
• How fast a product is formed (increases in mass or volume).

Mathematical Formula for Rate of Reaction

The average rate of reaction is calculated using the following mathematical formula:

Rate of reaction = Change in amount of reactant or productTime taken for change

Units of Rate of Reaction

The units depend on the state of the substances being measured and how we choose to measure them:

• If we measure the volume of a gas produced in cubic centimetres (cm³) over time in seconds (s), the unit of rate is cm³ s^-1 (cubic centimetres per second).
• If we measure the change in mass of a solid reactant in grams (g) over time in seconds (s), the unit of rate is g s^-1 (grams per second).
• If we measure the change in concentration of a solution in moles per cubic decimetre (mol dm^-3) over time in seconds (s), the unit of rate is mol dm^-3 s^-1 (moles per cubic decimetre per second).

Worked Example 1

During an experiment, a pupil reacted zinc granules with excess dilute hydrochloric acid. A volume of 48.0 cm³ of hydrogen gas was collected in exactly 120 seconds. Calculate the average rate of reaction.

Formula:

Average Rate = Volume of gas collectedTime taken

Substitution:

Average Rate = 48.0 cm³120 s

Answer:

Average Rate = 0.4 cm³ s^-1

Worked Example 2

In another experiment, the mass of a conical flask containing calcium carbonate and dilute nitric acid was measured. The initial mass was 250.60 g. After 60 seconds of reaction, the carbon dioxide gas escaped, and the new mass of the flask was 249.40 g. Calculate the rate of the reaction in g s^-1.

Step 1: Find the change in mass (mass of gas lost)

Mass lost = 250.60 g - 249.40 g = 1.20 g

Step 2: Calculate the rate

Rate = Loss in massTime taken

Rate = 1.20 g60 s

Rate = 0.02 g s^-1

FACTORS THAT AFFECT THE RATES OF CHEMICAL REACTIONS

To understand how different factors affect reaction rates, we use the Collision Theory.

The Collision Theory

The Collision Theory states that for a chemical reaction to occur:

• The reactant particles must collide with each other.
• The colliding particles must have a minimum amount of energy, known as the activation energy (E_a). Collisions with enough energy are called successful (or effective) collisions.
• The particles must collide in the correct orientation.

Any factor that increases the frequency of collisions or increases the energy of the colliding particles will increase the rate of the chemical reaction.

There are six main factors that affect the rate of a chemical reaction:

1. Temperature

Increasing the temperature of a reaction mixture increases the rate of reaction. When temperature increases, the reactant particles gain kinetic energy and move faster. This leads to two effects:

• The particles collide more frequently.
• The collisions are more energetic. A much larger fraction of the colliding particles now possess energy equal to or greater than the activation energy (E_a), leading to more successful collisions per unit time.

Example: Food spoils much faster at room temperature than when stored in a cold refrigerator because the chemical reactions causing decay slow down at lower temperatures. In the laboratory, magnesium reacts much faster with warm hydrochloric acid than with cold hydrochloric acid.

2. Concentration

Increasing the concentration of reactants in solution increases the rate of reaction. A higher concentration means there are more reactant particles packed into the same volume. This crowded environment increases the frequency of collisions between the particles, which increases the rate of successful collisions per unit time.

Example: Magnesium ribbon reacts slowly in dilute (0.5 mol dm^-3) hydrochloric acid, producing slow bubbles of hydrogen. When placed in concentrated (2.0 mol dm^-3) hydrochloric acid, the reaction is violent and fizzes rapidly.

3. Surface Area (Particle Size)

When one of the reactants is a solid, the reaction can only take place at the surface of the solid. Breaking a large solid reactant into smaller pieces (or a fine powder) increases its total surface area. This exposes more particles to the surrounding fluid reactant, increasing the frequency of collisions and thus increasing the rate of reaction.

Example: If you add a 5 g marble chip (calcium carbonate, CaCO₃) to dilute hydrochloric acid, it reacts slowly. If you grind the same 5 g marble chip into a fine powder and add it to the acid, it reacts extremely fast and fizzes vigorously.

4. Catalyst

A catalyst is a substance that increases the rate of a chemical reaction but remains chemically unchanged at the end of the reaction. A catalyst works by providing an alternative reaction pathway that has a lower activation energy (E_a). Because the barrier is lower, more reactant particles have the necessary energy to collide successfully, which increases the rate of reaction.

Example: Hydrogen peroxide (H₂O₂) decomposes very slowly at room temperature to form water and oxygen. If a small amount of manganese(IV) oxide (MnO₂) powder is added as a catalyst, the reaction happens instantly with rapid effervescence.

Word Equation: Hydrogen peroxide → Water + Oxygen

Chemical Equation: 2H₂O_2(aq) MnO₂ catalyst  2H₂O_(l) + O_2(g)

5. Pressure

This factor only affects reactions involving gases. Increasing the pressure of gaseous reactants pushes the gas particles closer together, decreasing the volume they occupy. This is equivalent to increasing the concentration of the gas. The particles collide more frequently, which increases the rate of reaction.

Example: In the industrial Haber process for manufacturing ammonia, a high pressure of about 200 atmospheres is maintained to speed up the reaction between nitrogen gas and hydrogen gas.

Chemical Equation: N_2(g) + 3H_2(g) ⇌ 2NH_3(g)

6. Light

Some reactions, called photochemical reactions, are initiated or accelerated by light energy (usually ultraviolet or visible light). Light particles (photons) provide the necessary activation energy to break chemical bonds in the reactant molecules, starting the reaction.

Example 1: Photosynthesis in green plants cannot take place in the dark. It requires sunlight to convert carbon dioxide and water into glucose and oxygen.

Example 2: Silver chloride (AgCl) is a white solid that decomposes into grey silver metal and chlorine gas when exposed to light. This reaction was the foundation of black-and-white photography.

Chemical Equation: 2AgCl_(s) sunlight  2Ag_(s) + Cl_2(g)

INTERPRET DATA ON THE RATE OF CHEMICAL REACTIONS

We can monitor how a reaction progresses over time by recording experimental data and plotting it on a graph. Analysing these graphs allows us to interpret the rate of chemical reactions at different stages.

Analysing a Typical Reaction Rate Graph

In a typical reaction where a gas is produced, we plot the Volume of gas produced on the vertical axis (y-axis) against Time on the horizontal axis (x-axis).

When we look at this curve, we can divide it into three distinct phases:

1. At the start (steepest curve): The slope of the curve is very steep. This tells us that the rate of reaction is at its highest. The reaction is fastest at the start because the concentration of reactant particles is at its maximum, leading to the highest frequency of successful collisions.
2. As the reaction proceeds (curve flattens): The slope of the curve becomes less steep, showing that the reaction is slowing down. This happens because reactant particles are being used up, so their concentration decreases. Collisions become less frequent.
3. At the end (flat horizontal line): The curve eventually levels off to a horizontal line. The slope is now zero, indicating that the reaction has stopped. This happens because at least one of the reactants (the limiting reactant) has been completely used up.

Worked Example of Data Interpretation

The table below shows the volume of carbon dioxide gas collected at 10-second intervals during the reaction of excess dilute hydrochloric acid with 2.0 g of calcium carbonate chips:

Questions and Worked Solutions:

Question A: At what time did the chemical reaction come to an end?

Answer: The reaction stopped at 50 seconds. This is because the volume of gas remained constant at 42.0 cm³ from 50 seconds onwards, showing that no more carbon dioxide was being produced.

Question B: Calculate the average rate of reaction during the first 30 seconds.

Working:

At time = 0 s, Volume = 0 cm³

At time = 30 s, Volume = 34.0 cm³

Change in volume = 34.0 cm³ - 0 cm³ = 34.0 cm³

Time taken = 30 s - 0 s = 30 s

Rate = Change in volumeTime taken

Rate = 34.0 cm³30 s

Answer: 1.13 cm³ s^-1

Question C: Compare the rate of reaction in the first 10 seconds to the rate of reaction between 30 and 40 seconds.

Working:

Rate in first 10 seconds: 15.0 cm³ - 0 cm³10 s = 1.50 cm³ s^-1

Rate between 30 and 40 seconds: 39.0 cm³ - 34.0 cm³10 s = 5.0 cm³10 s = 0.50 cm³ s^-1

Answer: The rate in the first 10 seconds (1.50 cm³ s^-1) is much faster than the rate between 30 and 40 seconds (0.50 cm³ s^-1) because the concentration of reactants was highest at the start of the reaction.

SUMMARY

• The rate of reaction measures how quickly reactants are converted into products over time.
• According to the Collision Theory, reactants must collide with sufficient energy (equal to or greater than the activation energy) and correct orientation to react.
• The rate of reaction is increased by:
  • Increasing the temperature: particles gain kinetic energy, moving faster and colliding more frequently with greater energy.
  • Increasing the concentration (or pressure of gases): places more particles in the same space, increasing collision frequency.
  • Increasing the surface area (smaller particle size of solids): exposes more reactant surface to collisions.
  • Adding a catalyst: provides an alternative pathway with a lower activation energy barrier.
  • Exposure to light: provides light energy to initiate photochemical reactions.
• Reaction progress graphs are steepest at the start (fastest rate), curve gently as reactants are consumed (slowing rate), and eventually flatten out to a horizontal line (reaction complete).

KEY TERMS

REVISION QUESTIONS

1. Define the term "rate of reaction" and state two standard units used to express it.

2. State the Collision Theory and explain what is meant by a "successful collision".

3. Explain why the rate of reaction between marble chips (calcium carbonate) and dilute hydrochloric acid is fastest at the very beginning of the reaction.

4. State and explain the effect of breaking a solid reactant into smaller pieces on the rate of reaction.

5. How does a catalyst speed up a chemical reaction? Draw a reference to activation energy in your explanation.

6. Explain why food stored in a freezer remains fresh for a longer period than food left on a kitchen counter in Zambia during summer.

7. Suggest two industrial reactions where catalysts are used to speed up operations.

8. Silver chloride decomposes when exposed to light. Write a balanced chemical equation for this reaction including state symbols.

PRACTICE EXERCISE

1. A student reacted a 3.0 g strip of magnesium ribbon with excess dilute sulphuric acid. The reaction finished in 40 seconds. Calculate the average rate of reaction in grams of magnesium consumed per second (g s^-1).

2. In an experiment, 50.0 cm³ of oxygen gas was collected in 2.5 minutes during the catalytic decomposition of hydrogen peroxide. Calculate the average rate of reaction in cm³ s^-1.

3. Dilute hydrochloric acid was added to excess sodium thiosulphate in a conical flask placed over a paper marked with a black cross. The cross became invisible due to the formation of a yellow precipitate of sulphur. The experiment was repeated at four different temperatures. The times taken for the cross to disappear are shown in the table below:

a) Plot a mental or physical graph of Time (y-axis) against Temperature (x-axis) and describe the trend shown.

b) Explain, in terms of the collision theory, why temperature has this effect on the reaction time.

c) Predict the time taken for the cross to disappear if the reaction was conducted at 45 °C.

4. Identify which of the following reactions will have a higher rate. Give a reason for your choice in each case:

a) Reaction A: Zinc powder and 1.0 mol dm^-3 HCl at 25 °C OR Reaction B: Large zinc granules and 1.0 mol dm^-3 HCl at 25 °C.

b) Reaction C: Magnesium ribbon and 2.0 mol dm^-3 H₂SO₄ at 30 °C OR Reaction D: Magnesium ribbon and 0.5 mol dm^-3 H₂SO₄ at 30 °C.

5. Gaseous sulphur dioxide reacts with oxygen to form sulphur trioxide in the Contact Process:

2SO_2(g) + O_2(g) ⇌ 2SO_3(g)

State two ways you could increase the rate of this gaseous reaction.

6. Copy and complete the table below by stating whether the rate of reaction will increase, decrease, or remain the same under the given changes:

ANSWERS TO PRACTICE EXERCISE

Answer 1:

Formula: Rate = Mass of magnesium consumedTime taken

Substitution: Rate = 3.0 g40 s

Working: 3.0 divided by 40 = 0.075

Answer: 0.075 g s^-1

Answer 2:

Convert time from minutes to seconds: 2.5 minutes = 2.5 × 60 s = 150 s

Formula: Rate = Volume of oxygen collectedTime taken

Substitution: Rate = 50.0 cm³150 s

Working: 50.0 divided by 150 = 0.3333...

Answer: 0.33 cm³ s^-1

Answer 3:

a) Trend: As the temperature increases, the time taken for the cross to disappear decreases. This shows that the rate of reaction increases with an increase in temperature.

b) Explanation: At higher temperatures, the thiosulphate and acid particles gain more kinetic energy and move faster. This causes more frequent collisions. Crucially, more particles have energy equal to or greater than the activation energy, increasing the frequency of successful collisions and thus speeding up the reaction.

c) Prediction: At 45 °C, the time taken will be approximately 15 seconds (midway between 20 seconds at 40 °C and 10 seconds at 50 °C).

Answer 4:

a) Reaction A will have a higher rate. Reason: Zinc powder has a much larger surface area than large granules, which increases the collision frequency between zinc atoms and acid molecules.

b) Reaction C will have a higher rate. Reason: 2.0 mol dm^-3 sulphuric acid has a higher concentration of hydrogen ions than 0.5 mol dm^-3 acid, leading to more frequent collisions per unit time.

Answer 5:

Any two of the following methods:

1. Increase the pressure of the reacting gases.
2. Increase the temperature of the reaction mixture.
3. Add a catalyst (such as vanadium(V) oxide, V₂O₅).

Answer 6: