# Dynamic Equilibrium ## Reversible Reactions - Some reactions go to completion where reactants are used up to form products - The reaction stops when all of the reactants are consumed - In **reversible reactions**, products can react to reform the original reactants - To show a reversible reaction, two opposing half arrows are used: $\rightleftharpoons$

A reversible reaction is one where products can react to reform the original reactants

- For example, the hydration and dehydration of copper(II) sulfate can be written as: - **Forward reaction**: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O (s)} \longrightarrow \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O (l)}$ - **Backward reaction**: $\text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O (l)} \longrightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O (s)}$ - **Combined equation**: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O (s)} \rightleftharpoons \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O (l)}$ ## What is Dynamic Equilibrium? - In a dynamic equilibrium, the reactants and products are **dynamic** (they are constantly moving)

A dynamic equilibrium exists when the rate of the forward reaction equals the rate of the backward reaction in a closed system, and the concentrations of reactants and products remain constant

### Key Features of Dynamic Equilibrium - **Both reactions continue**: The forward and backward reactions don't stop—they keep occurring - Think of it like an escalator with people going up at the same rate as people going down—everyone's moving but the number on each floor stays the same - **Equal rates**: The rate of the forward reaction = rate of the backward reaction - **Constant concentrations**: The concentrations of reactants and products don't change - However, these concentrations are **not necessarily equal** to each other - **Closed system required**: No reactants or products can escape ### Understanding Concentration Changes Over Time When equilibrium is approached **starting with reactants**: | Time Period | Product Concentration | Reactant Concentration | Status | |-------------|----------------------|------------------------|---------| | Initially | Increases rapidly | Decreases rapidly | Not at equilibrium | | As time progresses | Increase slows down | Decrease slows down | Approaching equilibrium | | At equilibrium | Remains constant (e.g., 10 mol dm$^{-3}$) | Remains constant (e.g., 1 mol dm$^{-3}$) | Equilibrium reached | - Note that the reactant concentration doesn't go to zero—there's still some left - The concentrations become constant at different values **Visual representation of concentration changes when starting with reactants:**

When equilibrium is approached **starting with products**: | Time Period | Reactant Concentration | Product Concentration | Status | |-------------|------------------------|----------------------|---------| | Initially | Increases from zero | Decreases | Not at equilibrium | | As time progresses | Increase slows down | Decrease slows down | Approaching equilibrium | | At equilibrium | Remains constant (e.g., 1 mol dm$^{-3}$) | Remains constant (e.g., 10 mol dm$^{-3}$) | Equilibrium reached | - **Important**: Whether you start with reactants or products, equilibrium concentrations end up at the same values - This demonstrates that equilibrium is independent of the starting conditions **Visual representation of concentration changes when starting with products:**

Notice how both charts reach the same final equilibrium concentrations (reactants = 1 mol dm$^{-3}$, products = 10 mol dm$^{-3}$), regardless of starting conditions. ### Visualizing Dynamic Equilibrium Consider the reaction: $\text{H}_2\text{(g)} + \text{I}_2\text{(g)} \rightleftharpoons 2\text{HI(g)}$ - At equilibrium, molecules of hydrogen iodide break down to form hydrogen and iodine at the same rate as hydrogen and iodine molecules combine to form hydrogen iodide - At any given moment, there are multiple molecules of each type present in the mixture - The exact composition stays constant even though individual molecules are constantly reacting

At equilibrium, concentrations are constant but not necessarily equal—don't confuse these two concepts

## Closed vs. Open Systems ### Closed Systems

A closed system is one in which none of the reactants or products can escape from the reaction mixture

- Equilibrium can only be achieved in a closed system - Matter cannot be lost to the surroundings - Energy can still be exchanged with surroundings (e.g., heat) **Example - Calcium Carbonate Decomposition in a Sealed Container**: $$\text{CaCO}_3\text{(s)} \rightleftharpoons \text{CaO(s)} + \text{CO}_2\text{(g)}$$ - In a sealed flask, no carbon dioxide can escape - The calcium carbonate, calcium oxide, and carbon dioxide reach equilibrium - Forward reaction: calcium carbonate decomposes to form calcium oxide and carbon dioxide - Backward reaction: calcium oxide reacts with carbon dioxide to reform calcium carbonate - At equilibrium, both reactions occur at the same rate ### Open Systems

An open system is one where matter and energy can be lost to the surroundings

**Special Cases Where Equilibrium Can Still Be Reached**: - When a reaction takes place **entirely in solution**, equilibrium can be reached in open flasks - This is because a negligible amount of material is lost through evaporation - Example: esterification reactions in aqueous solution - If the reaction involves **gases**, equilibrium can only be reached in a closed system - Gases readily escape if the container is open **Example - Calcium Carbonate Decomposition in an Open Container**: $$\text{CaCO}_3\text{(s)} \longrightarrow \text{CaO(s)} + \text{CO}_2\text{(g)}$$ - In an open beaker, carbon dioxide escapes continuously - The backward reaction cannot occur because there's no $\text{CO}_2$ present - The calcium carbonate continually decomposes until it's all gone - The reaction goes to completion rather than reaching equilibrium ### Summary Table: Closed vs. Open Systems | System Type | Can Equilibrium Be Reached? | Example | |------------|----------------------------|---------| | Closed (all states) | Yes | Sealed flask with $\text{CaCO}_3$ decomposition | | Open (gases involved) | No | Open beaker with $\text{CaCO}_3$ decomposition | | Open (only solutions, no significant gas loss) | Yes | Esterification in a beaker |

Reactions involving gases must occur in a closed system to reach equilibrium; otherwise, gaseous products will escape

## Common Misconceptions

The concentrations of reactants and products at equilibrium are NOT equal—they are constant (unchanging)

**Important clarifications**: - "Constant" means the values don't change once equilibrium is reached - "Equal" would mean the reactant and product concentrations have the same numerical value - At equilibrium, you might have 10 mol dm$^{-3}$ of products and 1 mol dm$^{-3}$ of reactants - These values are constant (they don't change) - But they are clearly not equal to each other - Before equilibrium is reached, concentrations **do change** as the reaction progresses - Only after equilibrium is established do the concentrations remain constant