Questions 10-13 refer to the following information. Cu(s) + 4 HNO3(aq) → Cu(NO3)2(aq) + 2 NO2(g) + 2 — Kinetics Chemistry Question

STEPS:

1. Understand the core chemistry concept (Collision Theory): For a chemical reaction to occur, reactant particles must collide with sufficient kinetic energy (equal to or greater than the activation energy, $E_a$) and in the correct spatial orientation. The rate of a chemical reaction is directly proportional to the frequency of these effective collisions.
2. Identify the factors that influence reaction rates: The rate of a reaction can be altered by changing parameters that affect collision frequency or collision energy:
* Concentration of aqueous/gaseous reactants (more particles per unit volume increases collision frequency).
* Temperature of the system (increases collision frequency and the fraction of particles with energy $\ge E_a$).
* Surface area of solid reactants (exposing more reactant particles to collisions).
* Catalysts (providing an alternative pathway with a lower $E_a$).
3. Analyze the effect of reducing reactant concentration:
* The concentration of the nitric acid solution used by the rest of the class was extremely high ($15.8\text{ M}$).
* If a student accidentally used a $1.5\text{ M}$ solution of $\text{HNO}_3(aq)$, the concentration of hydrogen ions and nitrate ions in the beaker would be reduced by more than a factor of 10.
* With significantly fewer reactant particles present per unit volume, the frequency of collisions between the aqueous acid ions and the solid copper metal surface decreases drastically, resulting in a much slower reaction rate.
4. Select the correct option: This identifies Option C as the only logical explanation for the slower rate.

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WHY_OTHERS_WRONG:

• Option A is incorrect: Smaller metal pieces have a larger total surface area-to-volume ratio. This exposes more copper atoms to the acid at any given moment, which increases collision frequency and results in a faster reaction rate, not a slower one.
• Option B is incorrect: Heating the reaction mixture increases the average kinetic energy of the particles. This increases both the collision frequency and the fraction of collisions that successfully overcome the activation energy barrier, leading to a faster reaction rate, not a slower one.
• Option D is incorrect: Using a larger mass of the solid mixture ($3.00\text{ g}$ instead of $2.00\text{ g}$) would increase the total amount of reactants and might increase the total surface area slightly, but it would not decrease the rate of collisions per unit of surface area. It would not cause the reaction to proceed at a slower rate.