Back to Chemistry syllabus

Chemical Reaction Rates and Equilibrium

CSCA Chemical Reaction Rates and Equilibrium study guide organized around the publicly available CSCA syllabus. Practice Chemistry questions on aicsca.com.

Before planning this topic, check the CSCA Exam Guide 2026 for exam dates, registration, fees, and subject requirements.

Syllabus Alignment

This study guide is organized around the publicly available CSCA syllabus for international undergraduate applicants.

Who It Is For

International students preparing for CSCA Math, Physics, Chemistry, or Chinese exams.

Related Practice

CSCA Practice · Go to questions

Past papers and worked video solutions · Timed mock exams · All subject video lessons

Practice by Topic

Jump from this tutorial to filtered practice questions for the same knowledge point.

Tutorial Content

Chemical Reaction Rates and Equilibrium

Core Concepts and Learning Objectives

This section studies chemical reactions from a dynamic perspective. It involves not only the change of substances (**Reaction Rate**) but also the limit of the reaction (**Chemical Equilibrium**).

**Learning Objectives:**

1. **Calculation**: Master the calculation of reaction rates and the relationship between rates of different substances.

2. **Mechanism**: Understand how concentration, temperature, pressure, and catalysts affect rates (Collision Theory).

3. **Equilibrium Criteria**: Deeply understand the "Dynamic, Equal, Constant, Changeable" characteristics and the Equilibrium Constant $K$.

4. **Dynamic Analysis**: Apply Le Chatelier's Principle to determine the direction of equilibrium shift.

---

I. Chemical Reaction Rate

#### 1. Definition and Expression

Measures how fast a reaction proceeds. Usually expressed as the decrease in reactant concentration or increase in product concentration per unit time.

* **Formula**: $v = \frac{|\Delta c|}{\Delta t}$

* **Unit**: $mol\cdot L^{-1}\cdot s^{-1}$ or $mol\cdot L^{-1}\cdot min^{-1}$

* **Key Rule**: For reaction $aA + bB \rightleftharpoons cC + dD$, the ratio of rates equals the ratio of stoichiometric coefficients:

$$ \frac{v(A)}{a} = \frac{v(B)}{b} = \frac{v(C)}{c} = \frac{v(D)}{d} $$

#### 2. Factors Affecting Rate

* **Intrinsic**: Nature of reactants (decisive factor).

* **Extrinsic** (Conditions):

1. **Concentration**: Higher concentration $\rightarrow$ More active molecules per unit volume $\rightarrow$ More frequent effective collisions $\rightarrow$ **Faster Rate**.

2. **Pressure** (Gases): Higher pressure $\equiv$ Smaller volume $\equiv$ Higher concentration $\rightarrow$ **Faster Rate**.

3. **Temperature**: Higher temperature $\rightarrow$ Percentage of activated molecules increases significantly $\rightarrow$ **Rate increases significantly**.

4. **Catalyst**: Lowers Activation Energy ($E_a$) $\rightarrow$ More molecules become activated $\rightarrow$ **Increases both forward and reverse rates equally**.

5. **Surface Area** (Solids): Crushing solids $\rightarrow$ Larger area $\rightarrow$ **Faster Rate**.

---

II. Chemical Equilibrium

#### 1. Establishment and Characteristics

Under given conditions, when **Forward Rate = Reverse Rate**, the concentrations of all substances stop changing. This is Chemical Equilibrium.

* **Characteristics**:

* **Reversible**: Applies to reversible reactions.

* **Dynamic**: $v_{fwd} = v_{rev} \neq 0$. Reaction hasn't stopped.

* **Equal**: Rates are equal.

* **Constant**: Concentrations remain constant.

* **Shift**: Equilibrium shifts if conditions change.

#### 2. Equilibrium Constant ($K$)

For $aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g)$, $K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}$.

* **Meaning**: Larger $K$ means the reaction proceeds further to the right.

* **Note**:

* $K$ depends **ONLY on Temperature**.

* Solids and pure liquids are **omitted** from the $K$ expression.

---

III. Le Chatelier's Principle

**Principle**: If a condition (concentration, pressure, temperature) acting on a system at equilibrium is changed, the system will shift to **counteract** the change.

| Change | Shift Direction | Reason |

| :--- | :--- | :--- |

| **Increase** Reactant Conc. | **Forward** (Right) | Consume added reactant |

| **Decrease** Product Conc. | **Forward** (Right) | Replace lost product |

| **Increase** Pressure (Volume $\downarrow$) | To side with **fewer** gas moles | Reduce pressure buildup |

| **Increase** Temperature | To **Endothermic** direction | Absorb excess heat |

| Add **Catalyst** | **No Shift** | Changes $v_{fwd}$ and $v_{rev}$ equally |

---

IV. Typical Examples

**Ex 1: Rate Conversion**

Reaction $N_2 + 3H_2 \rightleftharpoons 2NH_3$. If $v(NH_3) = 0.5$, what is $v(H_2)$?

* **Sol**: Ratio is $3:2$. $v(H_2) = \frac{3}{2} v(NH_3) = 0.75$.

**Ex 2: Pressure Effect**

$2NO_2(g) \rightleftharpoons N_2O_4(g)$ (Red-brown $\leftrightarrow$ Colorless). Compress the syringe to half volume.

* **Sol**:

1. **Instant**: Volume halves, concentration doubles $\rightarrow$ Color **darkens instantly**.

2. **Shift**: Pressure up, shifts to fewer moles (Right/Colorless).

3. **Final**: Color fades slightly from the peak but remains **darker** than the original state.

**Ex 3: Inert Gas (Trap)**

Add Helium (He) at constant temperature:

* **Case A (Constant Volume)**: Total P up, but partial pressures unchanged $\rightarrow$ **No Shift**.

* **Case B (Constant Pressure)**: Volume expands $\rightarrow$ Partial pressures drop $\rightarrow$ Shifts to side with **more** gas moles.

---

V. Common Pitfalls

1. **Start vs. Limit**: Catalysts help you reach equilibrium faster, but do not change the yield ($K$).

2. **Pressure Trap**: If gas moles are equal on both sides (e.g., $H_2 + I_2 \rightleftharpoons 2HI$), changing pressure changes Rate but **NOT Equilibrium**.

3. **Calculation Method**: Always use the **"ICE Table" (Initial, Change, Equilibrium)** method for calculations to avoid errors.