Chapter 3 / 3.2 Factors Influencing Rate of a Reaction

Why do some reactions zip along while others crawl? ⚡🐢

The speed (rate) of a reaction depends on how often reacting particles bump into each other with enough energy and the right orientation. Four ideas help us predict and control that speed. :contentReference[oaicite:0]{index=0}

1️⃣ Concentration – the starting line-up

Imagine players on a field: the more players, the more collisions. For a general reaction

\( aA + bB \;\longrightarrow\; cC + dD \)

experiments show the rate follows a rate law:

\( \text{Rate} = k\,[A]^x[B]^y \)

\(k\) is the rate constant (it changes only with temperature or catalyst).
\(x\) and \(y\) tell how strongly the rate depends on each reactant’s concentration. We find them only by experiment, not by the balanced equation. :contentReference[oaicite:1]{index=1}

Example (brown gas formation): Doubling \([NO]\) while keeping \([O_2]\) fixed makes the rate four times faster, so the law is

\( \text{Rate}=k\,[NO]^2[O_2] \)

This shows rate ∝ \([NO]^2\) but only ∝ \([O_2]\). 😎 :contentReference[oaicite:2]{index=2}

2️⃣ Order of reaction – the “power” game 💪

The order is the sum of the exponents in the rate law:

Order = \(x + y\)

So \( \text{Rate}=k[A]^{1/2}[B]^{3/2} \) is second-order, while \( \text{Rate}=k[A]^{3/2}[B]^{-1} \) is half-order! Orders can be 0, 1, 2, fractions… anything the experiment says. :contentReference[oaicite:3]{index=3}

Units of the rate constant help check the order:

Order	Units of \(k\)
0	mol L⁻¹ s⁻¹
1	s⁻¹
2	L mol⁻¹ s⁻¹

Spotting units like L mol⁻¹ s⁻¹ instantly tells you it’s a second-order reaction. 👌 :contentReference[oaicite:4]{index=4}

3️⃣ Molecularity – counting the collision squad 🤝

Molecularity is the number of particles that collide in a single elementary step:

Unimolecular: one particle decomposes, e.g., \( \mathrm{NH_4NO_2 \rightarrow N_2 + 2H_2O} \)
Bimolecular: two collide, e.g., \( 2HI \rightarrow H_2 + I_2 \)
Termolecular: three collide, e.g., \( 2NO + O_2 \rightarrow 2NO_2 \) (rare)

Key differences:

Order comes from experiments; molecularity is just a head-count in a single step.
Molecularity is always a whole number (never 0 or a fraction).
In complex (multi-step) reactions, only the slowest step decides the overall rate law. :contentReference[oaicite:5]{index=5}

4️⃣ Complex reactions & the slow-step 🐢

Think of a relay race: the slowest runner sets the team’s time. In chemistry, the rate-determining step does the same.

Example: In basic solution, iodide (\(I^-\)) speeds up hydrogen-peroxide breakdown:

\( \mathrm{H_2O_2 + I^- \rightarrow H_2O + IO^-} \) 🐢 (slow)
\( \mathrm{H_2O_2 + IO^- \rightarrow H_2O + I^- + O_2} \) (fast)

Because the first step is slow, the observed rate law is

\( \text{Rate}=k\,[H_2O_2][I^-] \)

The intermediate \(IO^-\) never shows up in the final equation but is crucial during the race! :contentReference[oaicite:6]{index=6}

🚀 Quick recap vs. NEET 🔥

Rate law & order: know how to get the rate law from data and sum exponents.
Units of \(k\): zero (mol L⁻¹ s⁻¹), first (s⁻¹), second (L mol⁻¹ s⁻¹).
Effect of concentration: changing \([A]\) by a factor \(n\) changes rate by \(n^x\).
Order vs. molecularity: experiment vs. collision count, whole vs. fractional.
Rate-determining step: for mechanisms, the slow step’s molecularity = overall order.

Master these five and you’ll breeze through typical kinetics questions on NEET! 🌟