Comprehensive Study Guide for Chemical Reaction Rates and Collision Theory

Fundamental Chemistry and Measurement Calculations

• Reactive Metals and Acids: When reactive metals are added to an acid, they produce hydrogen gas.     * Comparative Test: To determine whether magnesium or iron is more reactive, one could perform an experiment using water, acid, and the metal samples to observe the rate of hydrogen production.

• Definition of State Symbols:     * $(g)$: Gas     * $(l)$: Liquid     * $(s)$: Solid     * $(aq)$: Aqueous (in solution)

• Geometric Calculations for Cubes:     * Volume of a Large Cube ($3\,cm \times 3\,cm \times 3\,cm$): $27\,cm^2$     * Surface Area of a Large Cube ($3\,cm \times 3\,cm \times 3\,cm$): $54\,cm^3$     * Volume of 27 Smaller Cubes ($1\,cm \times 1\,cm \times 1\,cm$): (Not explicitly calculated in transcript logic beyond the 27 count).     * Surface Area of 27 Smaller Cubes ($1\,cm \times 1\,cm \times 1\,cm$): $162\,cm^3$     * Observation on Volume and Surface Area: The surface area was observed to be always larger than the volume in these specific examples.

Introduction to Chemical Kinetics and Collision Theory

• Reaction Velocity: Reactions vary in speed. Examples ordered from slowest to fastest:     1. Iron rusting (Slowest)     2. An egg frying     3. Fireworks exploding (Fastest)

• Collision Theory Requirements: A reaction occurs only when reacting particles collide with each other with enough energy. This is termed a successful collision.

• Activation Energy ($E_a$): This is the minimum amount of energy needed for colliding particles to react.

• Rate of Reaction Definition: The rate depends on the number of successful collisions that occur each second (the number of successful collisions per unit time).     * High Rate of Reaction: Associated with a large number of successful collisions per unit time.     * Low Rate of Reaction: Associated with a small number of successful collisions per unit time.

• Observation in Study: To study rates, observations must be chosen based on the state of the substance.     * Gases and solids are cited as the most useful for observation, especially if they are insoluble (do not dissolve).     * Liquids and solutions are less useful for visual observation changes as they typically appear similar to water.

Mathematical Determination of Reaction Rates

• Mean Rate of Reaction: This represents the average rate over the entire duration of the reaction.     * Formula: $\text{Mean rate of reaction} = \frac{\text{amount of product formed}}{\text{time taken to produce it}}$     * Example: If $40\,g$ of $CO_2$ is produced in $5\,s$     * $\text{Mean rate} = \frac{40\,g}{5\,s} = 8\,g/s$ (or $8\,gCO_2/s$)

• Instantaneous Rate of Reaction: This is the rate at a specific, particular time during the reaction.     * Calculation Method: The rate is equal to the gradient of the tangent line drawn to the curve on a graph of Mass of product ($g$) vs Time ($s$) at the specified time.     * Procedure:         1. Draw a tangent to the curve at the specified time.         2. Construct a triangle using the tangent as the hypotenuse/3rd side.         3. Measure the length of side "a" (vertical side in grams) and side "b" (horizontal side in seconds).         4. $\text{Gradient of tangent} = \frac{a}{b}$     * Worked Example: $80\,g$ in $5\,s$ results in an instantaneous rate of $16\,g/s$.

Investigation: Reaction of Marble Chips with Hydrochloric Acid

• Core Reaction Components: Pale grey marble chips are identified as impure calcium carbonate ($CaCO_3$). They react with a colourless solution of hydrochloric acid ($HCl$).

• Reaction Equation:     * Word: calcium carbonate + hydrochloric acid $\rightarrow$ calcium chloride + water + carbon dioxide     * Chemical: $CaCO_3 + HCl \rightarrow CaCl_2 + H_2O + CO_2$

• Measurement Rational: It is possible to measure the rate because a gas ($CO_2$) is produced, allowing for the measurement of the volume produced over a designated period of time.

• Experimental Apparatus: The setup involves a flask containing $5\,g$ of marble chips and $25\,ml$ of acid, connected to a gas syringe to capture the released gas.

Investigation: Effect of Concentration on Reaction Rate

• Reaction Scheme: Sodium thiosulfate + hydrochloric acid $\rightarrow$ sulfur + sulfur dioxide + sodium chloride + water.     * Chemical Equation: $Na_2S_2O_3(aq) + 2HCl(aq) \rightarrow S(s) + SO_2(g) + 2NaCl(aq) + H_2O(l)$

• Cloudiness Explanation: The solution becomes cloudy because sulfur is produced as a solid ($s$) in a very fine suspension. As more time passes, more solid is produced.

• Rate Measurement Technique: This reaction is measured by the "disappearing cross" method. The solution becomes a cloudy yellow, and the time taken for the solution to become opaque enough to hide a cross is measured.

• Procedural Necessity: It is essential to use different measuring cylinders for the acid and the thiosulfate to prevent the reaction from starting prematurely before the stop clock is activated.

• Experimental Results (Volume of Thiosulfate/Water/Time):     * $50\,cm^3$ Thiosulfate / $0\,cm^3$ Water: $156\,s$     * $40\,cm^3$ Thiosulfate / $10\,cm^3$ Water: $142\,s$     * $30\,cm^3$ Thiosulfate / $20\,cm^3$ Water: $192\,s$     * $20\,cm^3$ Thiosulfate / $130\,cm^3$ Water: $295\,s$ (Note: The horizontal data line for $10\,cm^3$ Thiosulfate is blank in the transcript).

• Analysis of Results:     * Most Concentrated: The first solution ($50\,cm^3$ thiosulfate : $0\,cm^{3}$ water).     * Slowest Reaction: The 20:30 solution ($295\,s$). (Note: Transcript mentions 20:30 but the table shows 20:130; discrepancy is preserved).     * Fastest Reaction: Theoretically, the 50:0 solution should be the fastest, but the student's results showed the 40:10 solution at $142\,s$ was faster, making the 50:0 result an anomaly.     * Variables:         * Independent: The ratio of thiosulfate and water.         * Dependent: Time taken for the cross to disappear.         * Controlled: Amount of hydrochloric acid, the specific cross used, and the total volume of solution ($50\,cm^3$).     * Trend: Higher volume of sodium thiosulfate correlates to a quicker time for the cross to disappear (e.g., $156\,s$ for $50\,cm^3$ vs $295\,s$ for $20\,cm^3$).     * Apparatus Change: If a boiling tube were used instead of a conical flask, the cross would disappear faster because the mixture would be deeper, requiring less opacity to obscure the cross as the surface area covering the cross is smaller.

Investigation: Effect of Surface Area (Particle Size)

• Scientific Context: Investigating surface area using the reaction between hydrochloric acid and marble chips.     * Reaction: $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$     * Mass Loss: The mass of the flask decreases over time because the carbon dioxide gas escapes.

• Experimental Setup:     * Apparatus: Conical flask containing acid and marble chips, placed on a balance.     * Cotton Wool: Placed in the neck of the flask to stop acid splashes from escaping while still allowing gas to leave.

• Methodology:     * Use the same mass ($6\,g$) of chips for each experiment to ensure a fair test by only changing the particle size.     * Measure $40\,cm^3$ of $HCl$.     * Record mass every $30\,s$ for at least 5 minutes.

• Findings:     * The calcium in the marble chips is the reacting species.     * Small vs. Large Chips: Small chips reacted more in the first 2 minutes than large chips, though they stopped reacting earlier. This demonstrates that a larger surface area leads to a quicker rate of reaction.

The Role of Temperature in Kinetic Energy

• General Effect: Increasing the temperature increases the rate of reaction.

• Equation: Uses the sodium thiosulfate reaction ($Na_2S_2O_3 + 2HCl$).

• Kinetics: Reacting particles gain energy and move faster at higher temperatures, increasing collision frequency. Collisions also have more energy, meaning more particles surpass the Activation Energy threshold.

Catalysts: Principles and Applications

• Definition: Catalysts are substances that speed up chemical reactions but are not used up; they remain chemically unchanged at the end.

• Decomposition of Hydrogen Peroxide:     * Equation: $2H_2O_2 \rightarrow 2H_2O + O_2$     * Catalyst: Manganese dioxide ($MnO_2$).     * Observation Data (Volume of $O_2$ vs Time):         * $0\,s$: $0\,cm^3$         * $10\,s$: $18\,cm^3$         * $20\,s$: $28\,cm^3$         * $30\,s$: $38\,cm^3$         * $40\,s$: $41\,cm^3$         * $50\,s$: $33\,cm^3$ (potential anomaly or reading error in transcript)         * $60\,s$: $42\,cm^3$         * $70\,s$: $43\,cm^3$         * $80\,s$: $42\,cm^3$         * $90\,s$: $44\,cm^3$         * $100\,s$: $44\,cm^3$         * $110\,s$: $44\,cm^3$     * Key Statistic: It took $14\,s$ for half of the oxygen to be produced.

• Mechanism of Action: Catalysts work by providing an "alberate route" for particles. This alternative pathway requires "tess Activation energy," meaning a greater fraction of particles possess sufficient energy to react.

• Energy Level Diagrams (Energy Profile):     * Reactants require energy to break existing chemical bonds.     * The catalyst lowers the Activation Energy requirement ($E_a$).     * By lowering this requirement, more successful collisions occur without needing to increase heat.

Advanced Collision Theory and Comprehensive Summaries

• Concentration Summary: Increasing concentration increases the number of reacting particles per $cm^3$. This increases the frequency of collisions and successful collisions, thereby increasing the rate.

• Surface Area Summary: Increasing surface area increases the number of particles able to collide. This increases collision frequency and successful collisions, raising the rate.

• Temperature Summary: Increasing temperature increases the energy of particles. They move faster and more particles collide with energy greater than the Activation Energy. Both energy and frequency of collisions increase.

• Thermal Decomposition of Potassium Chlorate:     * Definition: Thermal decomposition involves heat causing a chemical to break down.     * Reaction: $2KClO_3(s) \rightarrow 2KCl(s) + 3O_2(g)$     * Catalyst: Manganese (IV) oxide ($MnO_2$).     * Control Test: Heating $MnO_2$ alone produces no oxygen.     * Test for Oxygen: Relighting a glowing splint.     * Experimental Results:         1. Tube 1 ($KClO_3$ only): Produced oxygen second.         2. Tube 2 ($MnO_2$ only): No oxygen produced.         3. Tube 3 ($KClO_3 + MnO_2$): Produced oxygen first.

Practical Analysis and Revision Strategies

• Graph Interpretation:     * Curve A: Results for a mass of marble chips with $1.0\,mol\,dm^{-3}$ acid.     * Curve E: Represents the results if half the mass of marble chips is used, as half the chips produce half the volume of gas.     * Gradient: A steeper curve indicates a higher rate of reaction.

• Revision Tips from Mrs Ashton:     * Memorize experimental details and apparatus (e.g., purpose of cotton wool).     * Understand how to dilute reactants.     * Comment on variables and identify anomalies (and how to ignore/handle them).     * Practice drawing curves of best fit.     * Master Collision Theory, Activation Energy, and Particle Theory.     * Understand how catalysts work and how to prove they were not consumed.     * Knowledge of ingredients and reactants is essential for the 45-minute, 42-mark test.