Rates of Reaction and Reversible Reactions Vocabulary

Rates of Reaction

Introduction to Reaction Rates

• Definition: The rate of a chemical reaction is defined as how fast the reactants are transformed into products.
• Variation in Rates:
  • Slow reactions: An example is the rusting of iron.
  • Moderate speed reactions: An example is a metal, such as magnesium, reacting with acid to produce a gentle stream of bubbles.
  • Very fast reactions: An example is an explosion, which occurs almost instantaneously.

Factors Affecting the Rate of Reaction

There are four primary factors that determine the speed at which a chemical reaction occurs:

1. Temperature: Higher temperatures result in faster reactions.
2. Concentration (for solutions) or Pressure (for gases): More concentrated reactants or higher pressures lead to faster reaction rates.
3. Surface Area: This is determined by the size of solid reactant pieces. A larger surface area (achieved through smaller pieces or powders) results in a faster reaction.
4. Catalysts: The presence of a catalyst allows a reaction to proceed faster than it would without one.

Collision Theory and Activation Energy

• Collision Theory: This theory explains that the rate of a chemical reaction depends on two factors:
  1. Collision Frequency: How often the reacting particles collide. The more collisions that occur in a given amount of time, the faster the reaction proceeds.
  2. Energy Transfer: Particles must collide with sufficient energy for the collision to be "successful" (resulting in a reaction).
• Activation Energy: This is defined as the minimum amount of energy that particles must possess in order to react upon collision. If particles collide with energy lower than the activation energy, they simply bounce off one another without reacting.
• Ways to Increase Rate:
  • Increase the frequency of collisions to improve the probability of successful collisions.
  • Increase the energy of the collisions so a higher proportion of collisions result in a reaction.

Mechanisms of Increasing Reaction Rates

Temperature

• When temperature increases, particles move faster.
• Faster movement leads to more frequent collisions.
• Additionally, because particles move faster, they possess more energy. Therefore, a higher percentage of collisions will have energy exceeding the activation energy.

Concentration and Pressure

• Concentration: In a more concentrated solution, there are more reactant particles in the same volume of water. This makes collisions between reactant particles more likely.
• Pressure: In a gas, increasing the pressure squashes the particles closer together, increasing the frequency of collisions.

Surface Area

• If a reactant is solid, breaking it into smaller pieces increases its surface area to volume ratio.
• For the same volume of solid, more area is exposed for the particles in the solution to collide with.
• Powders offer the fastest reaction rates due to their exceptionally large surface area.

Rate and Proportionality

• The rate of a reaction is directly proportional to the frequency of successful collisions.
• For example, if the frequency of successful collisions doubles, the rate of reaction also doubles.

Catalysts and Enzymes

• Definition of a Catalyst: A substance that speeds up a reaction without being changed or used up in the process. Because they are not consumed, they do not appear in the chemical equation of the reaction (though they are sometimes written above the arrow).
• Mechanism of Action: Catalysts work by providing an alternative reaction pathway that has a lower activation energy ($E_a$).
• Enzymes: These are biological catalysts. They are specific, meaning an enzyme generally only catalyses one particular reaction for a specific molecule (e.g., enzymes in saliva breaking down starch or enzymes forming proteins).

Measuring Rates of Reaction

Calculating Mean Rate

• The mean reaction rate is calculated using the following formula:     $\text{Mean rate of reaction} = \frac{\text{Quantity of reactant used or product formed}}{\text{Time}}$
• Units of Rate:
  • If measuring mass over time: $\text{g/s}$
  • If measuring volume over time: $\text{cm}^3\text{/s}$
  • If measuring moles over time: $\text{mol/s}$
  • Time can also be measured in minutes, giving units like $\text{g/min}$.

Experimental Methods

1. Precipitation and Colour Change:
   • Used when the initial solution is transparent and the product is a solid precipitate that makes the solution "turbid" (cloudy).
   • Procedure: Observe a mark (like a black cross) through the solution and time how long it takes to disappear.
   • Disadvantage: The results are subjective as different people disagree on when the mark is truly gone. You cannot plot a rate graph using this single-point method.
2. Change in Mass:
   • Used for reactions that produce a gas. The reaction vessel is placed on a mass balance.
   • As gas is released, the mass disappears, and the balance reading drops.
   • A piece of cotton wool is placed in the neck of the flask to allow gas to escape while preventing reactant "spitting."
   • Advantage: Highly accurate and allows for plotting graphs over time.
   • Disadvantage: Releases gas into the room (dangerous if the gas is toxic).
3. Volume of Gas Given Off:
   • Uses a gas syringe to measure the volume of gas produced.
   • Advantage: Gas syringes are sensitive (usually to the nearest $\text{cm}^3$) and keep the gas contained.
   • Disadvantage: If the reaction is too vigorous, the plunger can be blown out of the syringe.

Rate of Reaction Graphs

• Slope and Gradient: The steepness of the curve represents the rate. A steeper line signifies a faster rate.
• Shape of the Curve: Graphs are usually curves that start steep (fastest rate due to high reactant concentration), get shallower (rate slows as reactants are used up), and eventually level off (reaction finished, one or more reactants used up).
• Comparing Rates:
  • The fastest reaction has the steepest initial slope and levels off earliest.
  • Reactions starting with the same amount of reactants will finish at the same level on the y-axis, even if their rates differ.

Calculating Instantaneous Rate Using Tangents

• To find the rate at a specific point in time, you calculate the gradient of a tangent to the curve at that point.
• Steps to Draw a Tangent:
  1. Place a ruler at the specific time point on the curve so it just touches it.
  2. Adjust until the space between the ruler and the curve is equal on both sides of the point.
  3. Draw the line.
• Calculating Gradient:$\text{Gradient} = \frac{\text{Change in } y}{\text{Change in } x}$
• Example calculation: For a mass/time graph at $40 \text{ s}$, if $\Delta y = 2.2 - 1.3$ and $\Delta x = 80 - 40$, then:     $\text{Rate} = \frac{0.90}{40} = 0.023 \text{ g/s}$

Required Practical 11: Investigating Reaction Rates

Magnesium and Hydrochloric Acid

• Equation: $Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)}$
• Procedure:
  1. Measure $50 \text{ cm}^3$ of dilute $HCl$ into a conical flask.
  2. Add magnesium ribbon and attach a gas syringe.
  3. Start the stopwatch and take volume readings at regular intervals (e.g., every $30 \text{ s}$) until the volume remains constant for three readings.
  4. Repeat with different acid concentrations while keeping the mass and surface area of magnesium, temperature, and volume of acid constant.

Sodium Thiosulfate and Hydrochloric Acid

• Equation: $Na_2S_2O_{3(aq)} + 2HCl_{(aq)} \rightarrow 2NaCl_{(aq)} + SO_{2(g)} + S_{(s)} + H_2O_{(l)}$
• Procedure:
  1. Add $50 \text{ cm}^3$ of dilute sodium thiosulfate to a flask.
  2. Place the flask on paper with a black cross.
  3. Add $10 \text{ cm}^3$ of dilute $HCl$ and start the stopwatch.
  4. Time how long it takes for the sulfur precipitate to make the cross disappear.
  5. Repeat with different concentrations. Use the same observer each time for fairness.

Reversible Reactions and Equilibrium

Nature of Reversible Reactions

• Definition: A reaction where the products can react together to reform the original reactants.
• Representation: Denoted by a double arrow: $A + B \rightleftharpoons C + D$.
• Condition Dependence: Changing conditions (temperature, pressure, concentration) can change the overall direction and the relative amounts of reactants and products.

Examples of Reversible Reactions

1. Ammonium Chloride:
   • Heating: $NH_4Cl_{(s)} \rightarrow NH_{3(g)} + HCl_{(g)}$ (Thermal decomposition).
   • Cooling: $NH_{3(g)} + HCl_{(g)} \rightarrow NH_4Cl_{(s)}$ (Recombination).
2. Hydrated Copper Sulfate:
   • $\text{Hydrated copper sulfate (blue)} \rightleftharpoons \text{Anhydrous copper sulfate (white)} + \text{Water}$
   • Heating blue crystals drives off water (endothermic) to leave white powder.
   • Adding water to the white powder reforms blue crystals (exothermic).
3. Haber Process:
   • $N_{2(g)} + 3H_{2(g)} \rightleftharpoons 2NH_{3(g)}$

Chemical Equilibrium

• Equilibrium State: Reached when the forward and reverse reactions occur at exactly the same rate.
• Closed System Required: Equilibrium can only be reached if no reactants or products can escape and nothing else can enter.
• Concentrations at Equilibrium: At equilibrium, the concentrations of reactants and products remain constant (they stop changing), but they are not necessarily equal.
  • Equilibrium lies to the right: Concentration of products is greater than reactants.
  • Equilibrium lies to the left: Concentration of reactants is greater than products.

Energy Transfer in Reversible Reactions

• If a reaction is endothermic in one direction, it must be exothermic in the opposite direction.
• Energy Conservation: The amount of energy absorbed by the endothermic reaction is exactly equal to the energy released by the exothermic reaction.

Le Chatelier's Principle

• The Principle: If the conditions of a reversible reaction at equilibrium are changed, the system will react to counteract that change.

Factors and Yield Predictions

1. Temperature:
   • Increase temperature: The system shifts to favor the endothermic reaction to absorb the extra energy. Yield of the endothermic direction increases.
   • Decrease temperature: The system shifts to favor the exothermic reaction to release more energy. Yield of the exothermic direction increases.
2. Pressure (Gases only):
   • Increase pressure: The system shifts toward the side with the smaller volume (the side with fewer gas molecules) to reduce pressure.
   • Decrease pressure: The system shifts toward the side with the larger volume (the side with more gas molecules) to increase pressure.
3. Concentration:
   • Increase reactant concentration: The system tries to decrease it by making more products (shifts right).
   • Decrease product concentration (removing product): The system tries to increase it again by converting more reactants into products (shifts right).

Science in Action: Industrial Compromise

• In industry, conditions are often a compromise. For example, a low temperature might increase the yield of an exothermic product (Le Chatelier), but it would also make the reaction rate too slow. A compromise temperature is chosen to get a reasonable yield at a reasonable speed.

Questions & Discussion

Retrieval Questions

• Q: What is the name of the principle which states that if you change conditions at equilibrium, the system tries to counteract the change?
  • A: Le Chatelier's Principle.
• Q: If you decrease the temperature, is the endothermic or exothermic reaction favoured?
  • A: Exothermic.
• Q: In a reversible reaction involving gases, what effect does increasing pressure have on the position of equilibrium?
  • A: It shifts the position of equilibrium to the side with fewer gas molecules.
• Q: What happens to the position of equilibrium if you decrease the concentration of products?
  • A: The equilibrium shifts to the right (toward the products) to replace what was lost.
• Q: What is a closed system?
  • A: A system where none of the reactants or products can escape and nothing else can get in.

Application Exercises

• Scenario 1: Ammonium Chloride Decomposition ($338 \,^{\circ}C$)
  • Equation: $NH_4Cl_{(s)} \rightleftharpoons NH_{3(g)} + HCl_{(g)}$ (Forward is endothermic).
  • i) If temperature is lowered to $250 \,^{\circ}C$, which reaction is favoured? A: The exothermic reverse reaction.
  • ii) Which compound will have a higher yield? A: Ammonium chloride ($NH_4Cl$).
  • iii) If temperature is raised to $400 \,^{\circ}C$, which reaction is favoured? A: The endothermic forward reaction.
  • iv) Which compounds will have a higher yield? A: Ammonia ($NH_3$) and hydrogen chloride ($HCl$).
• Scenario 2: Sulfur Trioxide Production
  • Equation: $2SO_{2(g)} + O_{2(g)} \rightleftharpoons 2SO_{3(g)}$
  • a) Which side has more volume? A: The left side (3 molecules of gas vs 2).
  • b) Effect of increasing pressure on $SO_3$ yield? A: Increases yield (shifts to the side with fewer molecules/less volume).
  • c) Effect of decreasing pressure on $SO_3$ yield? A: Decreases yield (shifts left).
• Scenario 3: Nitrogen Dioxide and Dinitrogen Tetroxide
  • Reaction: $2NO_{2(g)} \text{ (brown)} \rightleftharpoons N_2O_{4(g)} \text{ (colourless)}$
  • Observation: When compressed, the gas gets lighter in colour.
  • Explanation: Compression increases pressure. According to Le Chatelier’s principle, the system counteracts this by shifting to the side with fewer gas molecules. The left side has 2 molecules, the right has 1. Thus, it shifts right, producing more colourless $N_2O_4$ and making the mixture lighter.