Rate of Chemical Reaction 🚀

Chemistry tracks how quickly molecules change. Some reactions finish in a blink (mix AgNO3 with NaCl and boom—white AgCl forms!), while others crawl along for years (think of iron rusting). You’ll master both the numbers and the big ideas here. Let’s dive in! :contentReference[oaicite:0]{index=0}

1. What do we mean by “rate”? ⏱️

We measure speed as a change per unit time. For reactions, that change is concentration (mol L−1).

  • Disappearance of a reactant: \( \text{Rate} = -\dfrac{\Delta [R]}{\Delta t} \) :contentReference[oaicite:1]{index=1}
  • Appearance of a product: \( \text{Rate} = +\dfrac{\Delta [P]}{\Delta t} \) :contentReference[oaicite:2]{index=2}

Negative sign for reactants flips the value positive. Both expressions give the same rate, so use whichever species is easier to monitor! :contentReference[oaicite:3]{index=3}

2. Average vs. Instantaneous Rate 🎯

  1. Average rate looks at a finite time window. Example: hydrolyzing butyl chloride (C4H9Cl) drops from 0.100 M to 0.0905 M in 50 s, so \( r_{\text{avg}} = \dfrac{0.0905 – 0.100}{50} = 1.9 \times 10^{-4}\,\text{mol L}^{-1}\text{s}^{-1}\). :contentReference[oaicite:4]{index=4}
  2. Instantaneous rate zooms in to an infinitesimal moment: \( r_{\text{inst}} = -\dfrac{d[R]}{dt} = \dfrac{d[P]}{dt} \). Draw a tangent to the concentration-time curve and read its slope. The butyl-chloride plot gives \( r_{\text{inst}} = 5.1 \times 10^{-5}\,\text{mol L}^{-1}\text{s}^{-1} \) at 600 s. :contentReference[oaicite:5]{index=5}

3. Units of Rate 📏

Rate carries “concentration per time.” Common units:

  • \(\text{mol L}^{-1}\text{s}^{-1}\) (solution reactions)
  • \(\text{atm s}^{-1}\) (if you use partial pressures for gases) :contentReference[oaicite:6]{index=6}

4. Stoichiometry matters! ⚖️

When coefficients differ, divide by them so every species reports the same rate.

  • Example: \(2\,\mathrm{HI}\rightarrow \mathrm{H_2}+\mathrm{I_2}\)
    \( \displaystyle \text{Rate}= -\frac{1}{2}\frac{d[\mathrm{HI}]}{dt} = \frac{d[\mathrm{H_2}]}{dt} = \frac{d[\mathrm{I_2}]}{dt}\) :contentReference[oaicite:7]{index=7}
  • Bigger equation: \(5\,\mathrm{Br^-} + \mathrm{BrO_3^-} + 6\,\mathrm{H^+} \rightarrow 3\,\mathrm{Br_2} + 3\,\mathrm{H_2O}\)
    \( \displaystyle \text{Rate}= -\frac{1}{5}\frac{d[\mathrm{Br^-}]}{dt} = -\frac{1}{6}\frac{d[\mathrm{H^+}]}{dt} = \frac{1}{3}\frac{d[\mathrm{Br_2}]}{dt}\) :contentReference[oaicite:8]{index=8}

5. Expressing rate for gases 🌬️

At constant temperature, concentration is proportional to partial pressure, so you can swap [A] with PA and keep the same formulas. :contentReference[oaicite:9]{index=9}

6. Factors that influence rate 🔍

Four experimental knobs speed things up or slow them down:

  • Concentration (or pressure for gases)
  • Temperature
  • Presence of a catalyst
  • Overall pressure (for gaseous systems)

We explore each effect in depth later, but always remember: more collisions with enough energy = faster reaction!

High-Yield Ideas for NEET 📚

  1. Apply \( \text{Rate} = -\Delta[R]/\Delta t = +\Delta[P]/\Delta t \) to compute average rates quickly. :contentReference[oaicite:10]{index=10}
  2. Draw a tangent to a concentration-time curve to grab the instantaneous rate formula \( r_{\text{inst}} = -d[R]/dt \). :contentReference[oaicite:11]{index=11}
  3. Always divide by stoichiometric coefficients so every species yields the same rate value. :contentReference[oaicite:12]{index=12}
  4. Remember the standard unit \(\text{mol L}^{-1}\text{s}^{-1}\) and switch to \(\text{atm s}^{-1}\) for gaseous pressures. :contentReference[oaicite:13]{index=13}
  5. Predict rate changes qualitatively: higher concentration, temperature, or catalyst → faster reaction; opposite changes slow it down. :contentReference[oaicite:14]{index=14}

You’ve got this—keep practicing, and reaction-rate questions will feel like clockwork! ⏲️✨