Rates of Reaction and Collision Theory

Vocabulary flashcards covering methods of measuring reaction rates, graphical analysis, factors affecting rate, collision theory, and catalysis.

Rate of Reaction

The speed at which a chemical process occurs, measured as either the amount of reactant used up or the amount of product formed per unit time: $\text{Rate} = \frac{\text{amount of reactant used}}{\text{time taken}}$ or $\text{Rate} = \frac{\text{amount of product made}}{\text{time taken}}$.

Mass Loss Method

A method to determine reaction rate by recording the loss of mass on a balance over time as a dense gas (such as $CO_2$, $M_r = 44$) escapes from the reaction flask.

Downward Displacement of Water

A gas collection technique where gas produced in a flask travels through a delivery tube and displaces water inside an inverted measuring cylinder to measure volume over time.

Gas Syringe

A piece of laboratory glassware attached to a reaction flask to collect and measure the volume of an evolving gas over time, particularly useful for low molar mass gases such as hydrogen ($M_r = 2$).

Precipitation Method (Disappearing Cross)

A qualitative method to measure reaction rate by timing how long it takes for a solid precipitate (such as sulfur) to turn a clear solution opaque enough to obscure a black cross drawn beneath the flask.

Moles per Unit Time

A unit of reaction rate ($mol/s$) calculated by dividing the measured mass rate by the molar mass: $\text{moles per unit time} = \frac{\text{mass per unit time}}{\text{molar mass of the substance}}$.

Initial Rate Graph (Product)

A plot of product amount against time that displays a straight line with a positive correlation starting from the origin, where the gradient equals the initial rate of reaction.

Initial Rate Graph (Reactant)

A plot of reactant amount against time that displays a straight line with a negative correlation starting from the y-axis as the reactant is consumed.

Gradient of a Tangent

The slope of a straight line drawn to touch a curved rate graph at a specific time point, calculated using $\text{gradient} = \frac{\text{change in product}}{\text{change in time}}$ to determine the instantaneous rate of reaction.

Mean Rate of Reaction

The average rate of a reaction determined by calculating the overall change in reactant or product between two selected points along a reaction curve divided by the elapsed time.

Concentration Effect on Reaction Rate

An increase in solute concentration places more reactant particles within a given volume, causing more frequent collisions and a higher number of successful collisions per second.

Pressure Effect on Reaction Rate

An increase in the pressure of reacting gases compresses the same number of particles into a smaller volume, increasing collision frequency and the number of successful collisions per second.

Temperature Effect on Reaction Rate

An increase in temperature imparts higher kinetic energy to reactant particles, increasing collision frequency and resulting in a higher proportion of collisions having energy equal to or exceeding the activation energy.

Temperature Rule of Thumb

An empirical guideline stating that for aqueous and gaseous systems, every $10\,^{\circ}\text{C}$ increase in temperature approximately doubles the rate of reaction.

Surface Area Effect on Particle Size

Subdividing a solid into smaller pieces increases its total surface area (e.g., cutting a $2\,cm^3$ cube into $8$ smaller cubes increases surface area from $24\,cm^2$ to $48\,cm^2$), exposing more particles to collide per second.

Collision Theory

A chemical theory stating that reactions occur only when reactant particles collide with sufficient energy (greater than or equal to the activation energy) and the correct orientation.

Activation Energy ($E_a$)

The minimum amount of kinetic energy colliding reactant particles must possess in order to react successfully and form products.

Successful Collision

A collision between reactant particles possessing energy greater than or equal to the activation energy ($E \ge E_a$), leading to the rearrangement of atoms and formation of products.

Unsuccessful Collision

A collision where particles possess kinetic energy below the activation energy ($E < E_a$), causing them to rebound off one another without undergoing a chemical change.

Catalyst

A substance that speeds up a chemical reaction without being consumed or chemically altered itself, by providing an alternative reaction pathway that has a lower activation energy ($E_a$).

Reaction Profile (Catalyzed Reaction)

An energy diagram showing that a catalyzed reaction follows an alternative pathway with a reduced activation energy barrier ($E_a$) compared to the uncatalyzed reaction, while reactant and product energies remain unchanged.

Haber Process Catalyst

Iron beads used industrially at $450\,^{\circ}\text{C}$ and $200\,\text{atm}$ to speed up the synthesis of ammonia: $N_{2\,(g)} + 3H_{2\,(g)} \rightleftharpoons 2NH_{3\,(g)}$.