Science Unit 08

Chapter 8: Overview and Learning Objectives

• The study of reaction rates involves understanding how quickly products are formed from reactants or how fast reactants are consumed.

• Learners must be able to describe the various factors that influence chemical reaction rates.

• Key objectives include learning how a catalyst increases reaction rates while remaining chemically unchanged at the end of the process.

• Theoretical understanding centers on Collision Theory, using it to explain the effects of temperature, concentration, and surface area.

• Practical skills involve investigating reaction rates through experimentation and evaluating different methodologies for investigation.

Factors Influencing the Speed of a Reaction

• Research across a wide range of reactions highlights several major influences on reaction rate:

• Concentration Effects: Bringing reacting molecules closer together through the concentration of reactant solutions or the surface area of solid reactants.

• Pressure: Specifically for reacting gases, where increased pressure effectively increases the concentration of the gas particles.

• Temperature: The specific thermal level at which the reaction is performed.

• Catalysts: The introduction of substances that facilitate the reaction without being consumed.

Surface Area and Particle Size

• Chemical reactions involving solids can only occur at the surface of the solid where reactant particles can collide.

• Fragmentation Effect: Breaking a solid into smaller pieces increases the total surface area exposed to other reactants, increasing the probability of collisions.

• Mathematical Example (Cube Partitioning):  - A large cube of marble with side lengths of $2\,cm$ has a surface area of $24\,cm^2$.  - If this cube is cut into eight smaller cubes with side lengths of $1\,cm$, the total surface area increases to $48\,cm^2$.

• Practical Example with Calcium Carbonate (Marble):  - Experiment involves placing equal masses of marble in three forms—large lumps (A), small chips (B), and powder (C)—into equal volumes of dilute hydrochloric acid ($HCl$).  - Beaker C (powdered marble) will react the fastest because it has the largest surface area to volume ratio, allowing the most collisions between the acid and the carbonate.

• Iron and Oxygen:  - In a large lump of iron, oxygen molecules cannot access the internal atoms.  - In powdered iron, oxygen molecules can collide with many more iron atoms simultaneously, resulting in a significantly faster reaction.

Analysis of Reaction Progress Over Time

• A chemical reaction is typically fastest at the start because the concentration of reactant particles is at its peak.

• As the reaction progresses, it slows down because reactants are being used up, leading to a smaller change in the amount of product over the same time interval.

• The reaction finishes when no more product is formed, indicated by a flat line on a graph of product amount versus time.

• Graphing Principles:  - Time is usually plotted on the x-axis, and the amount/volume of product on the y-axis.  - A steeper curve indicates a faster initial reaction rate.  - The "line of best fit" should be drawn as a smooth curve or a straight line; data points should never be connected with jagged line segments.

Concentration and Reaction Rate

• Increasing the concentration of reactants in a solution increases the rate of reaction. This is because there are more particles in the same volume, which increases the frequency of collisions.

• Experimental Study (Marble Chips and Acid):  - Reaction: Marble chips plus dilute hydrochloric acid producing carbon dioxide ($CO_2$) gas.  - Setup: Uses a gas syringe connected to a test tube containing marble chips and acid to measure the volume of gas produced over time.  - Observation C ($1\,mol/dm^3$ acid) vs. Observation D ($0.5\,mol/dm^3$ acid):   - Curve C is steeper than Curve D, proving the reaction with more concentrated acid is faster.   - Initially, the reaction in C is twice as fast as in D, suggesting that doubling concentration can double the initial rate.   - Both reactions eventually produce the same total volume of gas (if the marble chips are the limiting reactant and in equal amounts), although C reaches this total faster.

• Units of Concentration: Concentrations are typically stated in $mol/dm^3$ or $g/dm^3$.

Pressure in Gaseous Reactions

• Increasing the pressure of gases is functionally equivalent to increasing the concentration; particles are forced closer together.

• Low pressure results in infrequent collisions between different molecules.

• High pressure makes collisions much more frequent.

• Industrial Applications:  - Haber Process: The industrial manufacture of ammonia ($NH_3$) by reacting nitrogen ($N_2$) with hydrogen ($H_2$) using an iron catalyst.  - Contact Process: The industrial manufacture of sulfuric acid ($H_2SO_4$) using sulfur and air.

Temperature and Kinetic Energy

• Increasing temperature causes atoms and molecules to move faster (increased kinetic energy).

• Particles collide more often and with greater force/energy.

• Food Preservation Example:  - Food is stored in refrigerators (approx. $3\,^\circ C$ to $5\,^\circ C$) or freezers ($-18\,^\circ C$ as recommended by the UK Food Standards Agency).  - Lower temperatures slow down the rate of decay and oxidation by air, keeping food fresh longer.

• Reaction Rates:  - A reaction at $40\,^\circ C$ will proceed significantly faster than the same reaction at $30\,^\circ C$ due to higher collision frequency and energy.

Collision Theory of Reaction Rate

• Collision Theory Definition: A theory stating that a chemical reaction takes place only when reactant particles collide with sufficient energy to initiate the reaction.

• Successful Collision Requirements:  1. Orientation: Molecules must collide with the correct geometric alignment.  2. Activation Energy ($E_a$): Molecules must collide with enough energy to overcome the minimum energy barrier required to break existing chemical bonds.

• The overall rate of reaction is determined by the rate of successful collisions.

Activation Energy ($E_a$)

• Activation energy is the minimum amount of energy needed to begin breaking bonds in reactant molecules.

• Every unique reaction has its own specific value for activation energy.

• If colliding particles have combined energy less than $E_a$, they will simply bounce off each other without reacting.

Catalysts: Mechanism and Enthalpy

• Catalyst Definition: A substance that speeds up a reaction without being chemically changed or consumed by the end of the process.

• Mechanism: A catalyst provides an alternative reaction pathway with a lower activation energy ($E_a$).

• Effect: By lowering the energy barrier, a higher proportion of collisions in the mixture have enough energy to be successful, increasing the reaction rate.

• Physical Properties: Catalysts usually have a large surface area and are often used in solid form for gaseous or liquid reactions.

• Enthalpy Diagrams:  - The enthalpy of reactants and products remains the same regardless of catalyst use.  - The "peak" of the energy curve (the activation energy barrier) is lower for the catalysed path compared to the uncatalysed path.

Biological Catalysts: Enzymes

• Enzymes are biological catalysts that facilitate metabolic reactions.

• They are highly sensitive to physical conditions, particularly temperature and $pH$.

• Common Enzyme Examples:  - Amylase: Found in the mouth; breaks down carbohydrates/sugars.  - Protease (e.g., Pepsin): Found in the stomach; breaks down proteins into amino acids.  - Lipase: Found in various body locations; breaks down fats into fatty acids and glycerol.

Questions & Discussion

• Question: What happens to reaction rate when surface area increases?

• Answer: The rate increases because more of the solid reactant is exposed, providing more sites for collisions.

• Question: Why is perishable food kept in a refrigerator?

• Answer: To slow down the rate of chemical reactions responsible for decay and oxidation by decreasing the thermal energy of the particles.

• Question: Describe the effect of using zinc powder vs. large pieces of zinc in sulfuric acid.

• Answer: Zinc powder provides a faster reaction rate due to increased surface area.

• Question: What is the effect of using a lower concentration of acid?

• Answer: The reaction rate decreases because there are fewer acid particles per unit volume, leading to lower collision frequency.

• Question: How does a catalyst change the activation energy?

• Answer: It decreases the activation energy required for the reaction to proceed.

• Question: What happens to the rate when samarium reacts with hot water instead of cold water?

• Answer: The rate increases because particles have more kinetic energy and collide more frequently and effectively.