Factors Influencing Rate of Reaction — Revision Notes

⚡ 30-Second Revision

Concentration — Rate $propto$ [Reactants]$^n$. Higher concentration $ightarrow$ more collisions $ightarrow$ faster rate.
Temperature — Rate $propto k$. $k = A e^{-E_a/RT}$. Higher T $ightarrow$ more energetic collisions $ightarrow$ faster rate (exponential effect).
Catalyst — Lowers $E_a$, provides alternative pathway. Increases rate without being consumed. Does not change $Delta G$ or $K_{eq}$.
Surface Area — For solids, larger surface area $ightarrow$ more contact points $ightarrow$ faster rate.
Nature of Reactants — Bond strength, physical state, complexity influence inherent reactivity.
Arrhenius Equation — lnleft(\frac{k_2}{k_1}\right) = \frac{E_a}{R}left(\frac{1}{T_1} - \frac{1}{T_2}\right)

2-Minute Revision

The speed of a chemical reaction, or its rate, is influenced by five primary factors. Firstly, reactant concentration: increasing it means more molecules per volume, leading to more frequent collisions and thus a faster rate, as described by the Rate Law.

Secondly, temperature: a rise in temperature dramatically increases the rate. This is because molecules gain more kinetic energy, leading to more frequent collisions, but more importantly, a significantly larger fraction of molecules achieve the necessary activation energy ($E_a$), as explained by the Arrhenius Equation.

Thirdly, a catalyst accelerates a reaction by providing an alternative pathway with a lower activation energy, without being consumed. It speeds up both forward and reverse reactions equally, so it doesn't affect equilibrium.

Fourthly, for reactions involving solids, increasing the surface area exposes more reactant particles, increasing contact points and collision frequency. Finally, the nature of reactants itself plays a role; some substances are inherently more reactive due to bond strengths, physical state, or molecular complexity.

Remember, these factors primarily work by increasing the frequency or effectiveness of molecular collisions.

5-Minute Revision

To master the factors influencing reaction rates for NEET, focus on the 'why' behind each effect. The core idea is collision theory: reactions need molecules to collide with sufficient energy ($E_a$) and correct orientation.

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Concentration — More reactants means more collisions. The Rate Law ($ext{Rate} = k[A]^x[B]^y$) quantifies this. If a reaction is first order in A, doubling [A] doubles the rate. If second order, it quadruples. For gases, increasing pressure (by decreasing volume) increases concentration and thus rate. Adding an inert gas at constant volume does NOT affect reactant concentrations, so no rate change.

* *Mini-Example*: If $ext{Rate} = k[X]^2$, and $[X]$ is tripled, new rate = $k(3[X])^2 = 9k[X]^2$, so rate increases 9 times.

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Temperature — A $10^circ C$ rise often doubles the rate. The Arrhenius Equation ($k = A e^{-E_a/RT}$) explains this. Higher T means molecules move faster (more collisions), but crucially, a much larger *fraction* of molecules possess energy $ge E_a$. This exponential increase in effective collisions is the main reason. Remember to use Kelvin for T in calculations. A plot of $ln k$ vs. $1/T$ gives a straight line with slope $-E_a/R$.

* *Mini-Example*: If $E_a$ is high, the rate is very sensitive to temperature changes.

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Catalyst — Speeds up reactions by providing an alternative reaction mechanism with a **lower activation energy ($E_a$)**. It does not get consumed, does not change $Delta G$, and does not shift equilibrium (it speeds up both forward and reverse reactions equally). Visualize this on an energy profile diagram: the 'hill' becomes smaller.

* *Mini-Example*: Enzymes in our body are biological catalysts, making complex reactions happen rapidly at body temperature.

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Surface Area — Relevant for heterogeneous reactions involving solids. Grinding a solid increases its surface area, exposing more reactant particles. This increases the number of contact points for collisions, thus increasing the rate.

* *Mini-Example*: Powdered sugar dissolves faster than a sugar cube.

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Nature of Reactants — Inherent properties matter. Ionic reactions are often faster than covalent ones (less bond breaking). Simpler molecules react faster than complex ones. Physical state (gas > liquid > solid) also affects mobility and collision frequency.

Key Formulas: $ext{Rate} = k[A]^x[B]^y$, lnleft(\frac{k_2}{k_1}\right) = \frac{E_a}{R}left(\frac{1}{T_1} - \frac{1}{T_2}\right). Practice numerical problems with these.

Prelims Revision Notes

Factors Influencing Rate of Reaction: NEET Revision Notes

1. Concentration of Reactants:

Effect: — Increasing reactant concentration generally increases the reaction rate.
Reason: — More molecules per unit volume $ightarrow$ increased collision frequency $ightarrow$ increased effective collisions.
Quantification: — Rate Law: $ext{Rate} = k[A]^x[B]^y$.

* $x, y$: Order of reaction w.r.t. A and B (experimentally determined). * Overall order = $x+y$.

Gases: — Increasing partial pressure of gaseous reactants (by decreasing volume) increases concentration and rate. Adding inert gas at constant volume does NOT affect rate.

2. Temperature:

Effect: — Increasing temperature significantly increases reaction rate (often doubles/triples for every $10^circ C$ rise).
Reason:

* Increased kinetic energy of molecules $ightarrow$ increased collision frequency (minor effect). * Crucial: Exponential increase in the fraction of molecules possessing energy $ge$ activation energy ($E_a$) (major effect, explained by Boltzmann distribution).

Quantification: — Arrhenius Equation: $k = A e^{-E_a/RT}$.

* lnleft(\frac{k_2}{k_1}\right) = \frac{E_a}{R}left(\frac{1}{T_1} - \frac{1}{T_2}\right). * $T$ must be in Kelvin. $R = 8.314,\text{J K}^{-1}\text{mol}^{-1}$. * Plot of $ln k$ vs. $1/T$ is linear with slope $= -E_a/R$.

3. Catalyst:

Effect: — Increases reaction rate without being consumed.
Mechanism: — Provides an alternative reaction pathway with a **lower activation energy ($E_a$)**.
Key Properties:

* Does not initiate reactions. * Does not change $Delta G$ or equilibrium constant ($K_{eq}$). Speeds up both forward and reverse reactions equally. * Specific in action. * Effective in small amounts.

Types: — Homogeneous (same phase), Heterogeneous (different phase), Enzyme (biological).

4. Surface Area (for heterogeneous reactions with solids):

Effect: — Increasing surface area increases reaction rate.
Reason: — More exposed surface $ightarrow$ more contact points/active sites $ightarrow$ increased collision frequency between solid and other reactants.
Example: — Powdered zinc reacts faster with acid than a lump of zinc.

5. Nature of Reactants:

Effect: — Inherent chemical properties influence reactivity.
Factors:

* Bond Strength: Weaker bonds break more easily, leading to faster reactions. * Physical State: Gases and liquids react faster than solids due to higher molecular mobility. * Complexity: Simpler molecules/ions react faster than complex ones (less rearrangement needed). * Ionic vs. Covalent: Ionic reactions in solution are often very fast.

6. Presence of Radiation/Light:

Effect: — For photochemical reactions, light absorption can initiate or accelerate reactions by exciting molecules or forming radicals.

Common Misconceptions:

Catalysts initiate reactions or change equilibrium (False).
Order of reaction = stoichiometric coefficient (False, only for elementary reactions).
Temperature only increases collision frequency (False, mainly increases effective collisions).

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Concentration
Temperature
Catalyst
Surface Area
Nature of Reactants