Zero and First Order Reactions

Q: What is the primary difference between the rate constant units for zero and first-order reactions?

The units of the rate constant ($k$) are distinct for different reaction orders. For a zero-order reaction, the rate is independent of concentration, so the rate constant has units of concentration per unit time, typically mol L$^{-1}$ s$^{-1}$. This is because Rate = $k[A]^0 = k$. For a first-order reaction, the rate is directly proportional to the concentration, so the rate constant has units of inverse time, typically s$^{-1}$. This is derived from Rate = $k[A]$, so $k = \text{Rate}/[A] = (\text{mol L}^{-1} \text{s}^{-1}) / (\text{mol L}^{-1}) = \text{s}^{-1}$. Understanding these units is crucial for identifying reaction order from given rate constant values.

Q: Why is the half-life of a first-order reaction independent of the initial concentration?

The independence of half-life ($t_{1/2}$) from initial concentration for a first-order reaction is a unique and important characteristic. The integrated rate law for a first-order reaction is $\ln([A]_t/[A]_0) = -kt$. When $[A]_t = [A]_0/2$, we substitute this into the equation to get $t_{1/2} = \ln(2)/k = 0.693/k$. As you can see, the expression for $t_{1/2}$ only contains the rate constant $k$ (which is temperature-dependent) and a constant (0.693). It does not include $[A]_0$. This means that no matter how much reactant you start with, it will always take the same amount of time for half of it to react.

Updated 22 Mar 2026

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The order of a chemical reaction, with respect to a particular reactant, is defined as the exponent to which its concentration term is raised in the experimentally determined rate law. The overall order of a reaction is the sum of the exponents of the concentration terms in the rate law. Zero-order reactions are those where the rate of reaction is independent of the concentration of the reactant(s…

Quick Summary

Zero and first-order reactions are fundamental concepts in chemical kinetics, describing how reaction rates depend on reactant concentrations. A zero-order reaction proceeds at a constant rate, entirely independent of the reactant's concentration.

Its integrated rate law is $[A]_t = [A]_0 - kt$ , and a plot of $[A]_t$ vs. time yields a straight line with slope $-k$ . The half-life ( $t_{1/2} = [A]_0 / 2k$ ) is directly proportional to the initial concentration.

The rate constant $k$ has units of mol L $^{-1}$ s $^{-1}$ . Examples include enzyme-saturated reactions or surface-catalyzed reactions.

A first-order reaction has a rate directly proportional to the first power of the reactant's concentration. Its integrated rate law is $\ln([A]_t/[A]_0) = -kt$ (or $2.303 \log([A]_t/[A]_0) = -kt$ ), and a plot of $ln[A]_t$ vs.

time gives a straight line with slope $-k$ . Crucially, its half-life ( $t_{1/2} = 0.693/k$ ) is constant and independent of the initial concentration. The rate constant $k$ has units of s $^{-1}$ . Radioactive decay is a classic example.

Understanding these distinctions, including their integrated rate laws, half-life expressions, and graphical representations, is vital for NEET.

Full explanation

Chemical kinetics is the branch of chemistry that deals with the rates of chemical reactions and the factors influencing them. A fundamental concept in kinetics is the 'order of reaction', which describes how the rate of a reaction depends on the concentration of its reactants. It's an experimentally determined value, not necessarily derived from the stoichiometry of the balanced chemical equation.

Conceptual Foundation: Rate Law and Order of Reaction

The Rate Law expresses the relationship between the rate of a reaction and the concentrations of the reactants. For a general reaction $aA + bB \rightarrow cC + dD$ , the rate law is typically written as:

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

where:

\text{Rate}

is the speed at which reactants are consumed or products are formed.

k

is the rate constant, a proportionality constant specific to a given reaction at a particular temperature.

[A]

and

[B]

are the molar concentrations of reactants A and B.

x

and

y

are the orders of reaction with respect to reactants A and B, respectively. These are experimentally determined exponents and can be integers, fractions, or even zero.

The Overall Order of Reaction is the sum of the individual orders, i.e., $x+y$ . It's important to distinguish between the order of reaction and molecularity. Molecularity refers to the number of reacting species (atoms, ions, or molecules) that collide simultaneously in an elementary step of a reaction. It is always an integer and applies only to elementary reactions, whereas order can be for elementary or complex reactions and can be non-integer.

Key Principles and Laws: Zero-Order Reactions

A reaction is said to be zero-order if its rate is independent of the concentration of the reactant. This means the exponent of the reactant concentration in the rate law is zero.

Consider a general zero-order reaction: $A \rightarrow \text{Products}$

1. Differential Rate Law:

\text{Rate} = -\frac{d[A]}{dt} = k[A]^0 = k

Since

[A]^0 = 1

, the rate is simply equal to the rate constant

k

2. Integrated Rate Law:

To find how the concentration of A changes over time, we integrate the differential rate law:

-\frac{d[A]}{dt} = k

d[A] = -k \, dt

Integrating from initial concentration

[A]_0

at time

t=0

to concentration

[A]_t

at time

t

\int_{[A]_0}^{[A]_t} d[A] = -k \int_0^t dt

[A]_t - [A]_0 = -kt

[A]_t = [A]_0 - kt

This is the integrated rate law for a zero-order reaction. It shows a linear decrease in concentration over time.

3. Characteristics of Zero-Order Reactions:

Rate: — Constant and independent of reactant concentration.

Units of Rate Constant ($k$): — Since

\text{Rate} = k

and Rate has units of concentration/time (e.g., mol L

^{-1}

^{-1}

), the units of

k

for a zero-order reaction are also mol L

^{-1}

^{-1}

Graphical Representation: — A plot of

[A]_t

versus time (

t

) yields a straight line with a negative slope equal to

-k

and a y-intercept equal to

[A]_0

* Slope $= -k$ * Y-intercept $= [A]_0$

Half-life ($t_{1/2}$): — The time required for the concentration of a reactant to decrease to half its initial value. At

t = t_{1/2}

[A]_t = [A]_0 / 2

Substituting into the integrated rate law:

\frac{[A]_0}{2} = [A]_0 - k t_{1/2}

k t_{1/2} = [A]_0 - \frac{[A]_0}{2} = \frac{[A]_0}{2}

t_{1/2} = \frac{[A]_0}{2k}

For a zero-order reaction, the half-life is directly proportional to the initial concentration. This means it takes longer for half of a larger initial amount to react.

4. Real-World Applications:

Enzyme-catalyzed reactions often exhibit zero-order kinetics when the substrate concentration is much higher than the enzyme concentration, and the enzyme active sites are saturated. The rate is then limited by the enzyme's turnover rate, not the substrate amount.

Reactions occurring on a metal surface, like the decomposition of ammonia on a hot platinum surface (

2NH_3(g) \xrightarrow{Pt} N_2(g) + 3H_2(g)

), can be zero-order if the surface is fully covered by reactant molecules. The rate is then limited by the surface area, not the gas phase concentration.

Photochemical reactions where the rate is limited by the intensity of light absorbed, rather than the reactant concentration.

Key Principles and Laws: First-Order Reactions

A reaction is said to be first-order if its rate is directly proportional to the first power of the concentration of one reactant.

Consider a general first-order reaction: $A \rightarrow \text{Products}$

1. Differential Rate Law:

\text{Rate} = -\frac{d[A]}{dt} = k[A]^1 = k[A]

2. Integrated Rate Law:

To find how the concentration of A changes over time, we integrate the differential rate law:

-\frac{d[A]}{dt} = k[A]

\frac{d[A]}{[A]} = -k \, dt

Integrating from initial concentration

[A]_0

at time

t=0

to concentration

[A]_t

at time

t

\int_{[A]_0}^{[A]_t} \frac{d[A]}{[A]} = -k \int_0^t dt

[\ln[A]]_{[A]_0}^{[A]_t} = -k[t]_0^t

\ln[A]_t - \ln[A]_0 = -kt

\ln\left(\frac{[A]_t}{[A]_0}\right) = -kt

This is one form of the integrated rate law for a first-order reaction.

It can also be written as:

[A]_t = [A]_0 e^{-kt}

Or, converting natural logarithm to base-10 logarithm:

2.303 \log\left(\frac{[A]_t}{[A]_0}\right) = -kt

$$ \log\left(\frac{[A]_t}{[A]_0}\right) = -\frac{kt}{2.

303}

\log[A]_t = \log[A]_0 - \frac{kt}{2.

3. Characteristics of First-Order Reactions:

Rate: — Directly proportional to the first power of reactant concentration.

Units of Rate Constant ($k$): — Since

\text{Rate} = k[A]

, then

k = \frac{\text{Rate}}{[A]}

. Units of Rate are mol L

^{-1}

^{-1}

and units of

[A]

are mol L

^{-1}

. Therefore, units of

k

are

\frac{\text{mol L}^{-1} \text{s}^{-1}}{\text{mol L}^{-1}} = \text{s}^{-1}

Graphical Representation: — A plot of

ln[A]_t

versus time (

t

) yields a straight line with a negative slope equal to

-k

and a y-intercept equal to

ln[A]_0

. Similarly, a plot of

log[A]_t

versus time (

t

) yields a straight line with a negative slope equal to

-\frac{k}{2.303}

and a y-intercept equal to

log[A]_0

* Slope $= -k$ (for $ln[A]_t$ vs $t$ ) * Slope $= -k/2.303$ (for $log[A]_t$ vs $t$ )

Half-life ($t_{1/2}$): — At

t = t_{1/2}

[A]_t = [A]_0 / 2

Substituting into the integrated rate law $\ln\left(\frac{[A]_t}{[A]_0}\right) = -kt$ :

\ln\left(\frac{[A]_0/2}{[A]_0}\right) = -k t_{1/2}

\ln\left(\frac{1}{2}\right) = -k t_{1/2}

-\ln(2) = -k t_{1/2}

t_{1/2} = \frac{\ln(2)}{k} = \frac{0.693}{k}

For a first-order reaction, the half-life is constant and independent of the initial concentration of the reactant. This is a very important characteristic.

4. Real-World Applications:

Radioactive decay: — All radioactive decay processes follow first-order kinetics. For example, the decay of Carbon-14 used in radiocarbon dating.

Decomposition reactions: — Many unimolecular decomposition reactions in the gas phase, such as the decomposition of

N_2O_5

(

N_2O_5(g) \rightarrow N_2O_4(g) + \frac{1}{2}O_2(g)

), follow first-order kinetics.

Hydrolysis of esters in acidic medium: — While the overall reaction might seem second order, if water is in large excess (solvent), its concentration remains effectively constant, making it a pseudo-first-order reaction.

Common Misconceptions and NEET-Specific Angle

Order vs. Molecularity: — Students often confuse these. Remember, order is experimental and can be fractional or zero; molecularity is theoretical (for elementary steps) and always an integer (1, 2, or 3).

Units of Rate Constant: — The units of

k

depend on the order of the reaction. For zero-order, it's mol L

^{-1}

^{-1}

. For first-order, it's s

^{-1}

. This is a common MCQ question.

Half-life Dependence: — A critical distinction is the dependence of

t_{1/2}

on initial concentration. For zero-order,

t_{1/2} \propto [A]_0

. For first-order,

t_{1/2}

is independent of

[A]_0

. This is a frequent basis for numerical problems and conceptual questions.

Graphical Interpretation: — Be adept at interpreting plots of concentration vs. time,

\ln(\text{concentration})

vs. time, and

\log(\text{concentration})

vs. time to determine the order of a reaction and calculate the rate constant.

Integrated Rate Laws: — Memorize and understand the derivation of the integrated rate laws and half-life expressions for both zero and first-order reactions. NEET questions often involve direct application of these formulas or require calculating one parameter given others.

Pseudo-First-Order Reactions: — Understand that a higher-order reaction can behave as first-order if one reactant is in vast excess, effectively making its concentration constant. This simplifies the kinetics to first-order.

Mastering these concepts, derivations, and their applications is essential for tackling NEET questions on chemical kinetics.

Key Concepts

Integrated Rate Law for Zero-Order Reactions

The integrated rate law for a zero-order reaction, $A \rightarrow \text{Products}$ , is $[A]_t = [A]_0 - kt$ .…

Integrated Rate Law for First-Order Reactions

For a first-order reaction, $A \rightarrow \text{Products}$ , the integrated rate law is $\ln([A]_t/[A]_0) =…

Half-life (

t_{1/2}

) for Zero and First-Order Reactions

The half-life is a characteristic time for a reaction. For a zero-order reaction, $t_{1/2} = [A]_0 / 2k$ .…

Often confused with

Side-by-side differences the NEET paper likes to test.

Zero and First Order Reactions vs First-Order Reactions

Aspect	Zero and First Order Reactions	First-Order Reactions
Rate Law	Rate = $k[A]^0 = k$	Rate = $k[A]^1 = k[A]$
Integrated Rate Law	$[A]_t = [A]_0 - kt$	$\ln([A]_t/[A]_0) = -kt$ or $2.303 \log([A]_t/[A]_0) = -kt$
Units of Rate Constant ($k$)	Concentration/Time (e.g., mol L$^{-1}$ s$^{-1}$)	Time$^{-1}$ (e.g., s$^{-1}$)
Half-life ($t_{1/2}$)	$t_{1/2} = [A]_0 / 2k$ (depends on initial concentration)	$t_{1/2} = 0.693 / k$ (independent of initial concentration)
Graphical Plot for Linearity	$[A]_t$ vs. $t$ (slope = $-k$)	$\ln[A]_t$ vs. $t$ (slope = $-k$)
Effect of Doubling [A]	Rate remains unchanged	Rate doubles

Zero-order reactions have a constant rate, independent of reactant concentration, with a half-life directly proportional to the initial concentration. Their rate constant units are concentration per time.

In contrast, first-order reactions have a rate directly proportional to the reactant concentration, and their half-life is constant, independent of the initial concentration. Their rate constant units are inverse time.

These differences are critical for identifying reaction order and solving related numerical problems in NEET.

Why it is tested: NEET relevance: High. Understanding these distinctions is fundamental for solving numerical problems involving integrated rate laws, half-life calculations, and interpreting experimental data (especially graphical plots) to determine reaction order. These concepts are frequently tested.

Questions students ask

5 answered on this topic.

What is the primary difference between the rate constant units for zero and first-order reactions?

The units of the rate constant ( $k$ ) are distinct for different reaction orders. For a zero-order reaction, the rate is independent of concentration, so the rate constant has units of concentration per unit time, typically mol L $^{-1}$ s $^{-1}$ .

This is because Rate = $k[A]^0 = k$ . For a first-order reaction, the rate is directly proportional to the concentration, so the rate constant has units of inverse time, typically s $^{-1}$ . This is derived from Rate = $k[A]$ , so $k = \text{Rate}/[A] = (\text{mol L}^{-1} \text{s}^{-1}) / (\text{mol L}^{-1}) = \text{s}^{-1}$ .

Understanding these units is crucial for identifying reaction order from given rate constant values.

How can I graphically distinguish between a zero-order and a first-order reaction?

Graphical methods are excellent for determining reaction order. For a zero-order reaction, plotting the concentration of reactant $[A]_t$ against time ( $t$ ) will yield a straight line with a negative slope equal to $-k$ .

For a first-order reaction, plotting the natural logarithm of the reactant concentration, $ln[A]_t$ , against time ( $t$ ) will yield a straight line with a negative slope equal to $-k$ . Alternatively, plotting $log[A]_t$ vs $t$ for a first-order reaction also gives a straight line with a slope of $-k/2.

303$. If these specific plots are linear, they confirm the respective reaction order.

Why is the half-life of a first-order reaction independent of the initial concentration?

The independence of half-life ( $t_{1/2}$ ) from initial concentration for a first-order reaction is a unique and important characteristic. The integrated rate law for a first-order reaction is $\ln([A]_t/[A]_0) = -kt$ .

When $[A]_t = [A]_0/2$ , we substitute this into the equation to get $t_{1/2} = \ln(2)/k = 0.693/k$ . As you can see, the expression for $t_{1/2}$ only contains the rate constant $k$ (which is temperature-dependent) and a constant (0.

693). It does not include $[A]_0$ . This means that no matter how much reactant you start with, it will always take the same amount of time for half of it to react.

Can a reaction be both zero-order and first-order under different conditions?

Yes, it's possible for a reaction to exhibit different orders under varying conditions. For instance, many enzyme-catalyzed reactions follow Michaelis-Menten kinetics. At very low substrate concentrations, the reaction can be first-order with respect to the substrate.

However, at very high substrate concentrations, the enzyme active sites become saturated, and the reaction rate becomes independent of further increases in substrate concentration, thus behaving as a zero-order reaction.

Similarly, heterogeneous catalytic reactions can switch orders depending on surface coverage. The order is an experimental observation, not an inherent property of the reactants themselves.

What is a pseudo-first-order reaction, and why is it important for NEET?

A pseudo-first-order reaction is a higher-order reaction (typically second or third order) that behaves like a first-order reaction because the concentration of one or more reactants is kept in vast excess.

For example, the hydrolysis of an ester in the presence of a large excess of water (solvent) is actually a second-order reaction (first order with respect to ester, first order with respect to water).

However, since water's concentration remains virtually constant throughout the reaction, its concentration term gets absorbed into the rate constant, making the observed reaction first-order with respect to the ester.

This concept is important for NEET as it tests your understanding of how experimental conditions can simplify complex kinetics and how to identify the true order versus the observed order.

Revise in 30 seconds

Zero-Order Reaction:

- Rate Law: $\text{Rate} = k$ - Integrated Rate Law: $[A]_t = [A]_0 - kt$ - Half-life: $t_{1/2} = \frac{[A]_0}{2k}$ (proportional to $[A]_0$ ) - Units of $k$ : mol L $^{-1}$ s $^{-1}$ - Linear Plot: $[A]_t$ vs. $t$ (slope = $-k$ )

First-Order Reaction:

- Rate Law: $\text{Rate} = k[A]$ - Integrated Rate Law: $\ln\left(\frac{[A]_t}{[A]_0}\right) = -kt$ or $2.303 \log\left(\frac{[A]_0}{[A]_t}\right) = kt$ - Half-life: $t_{1/2} = \frac{0.693}{k}$ (independent of $[A]_0$ ) - Units of $k$ : s $^{-1}$ - Linear Plot: $\ln[A]_t$ vs. $t$ (slope = $-k$ )

Zero Constant Linear Half-life Proportional

Zero-order: Constant rate (independent of concentration)

Linear plot:

[A]

t

Half-life: Proportional to initial concentration (

[A]_0

)

First Exponential Log Half-life Independent

First-order: Exponential decay

Log plot:

\ln[A]

t

Half-life: Independent of initial concentration