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HSC Chemistry

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  1. Module 1: Properties and Structure of Matter
    1.1 Properties of Matter
  2. 1.2 Atomic Structure and Atomic Mass
  3. 1.3 Periodicity
  4. 1.4 Bonding
  5. Module 2: Introduction to Quantitative Chemistry
    2.1 Chemical Reactions and Stoichiometry
  6. 2.2 Mole Concept
  7. 2.3 Concentration and Molarity
  8. 2.4 Gas Laws
  9. Module 3: Reactive Chemistry
    3.1 Chemical Reactions
  10. 3.2 Predicting Reactions of Metals
  11. 3.3 Rates of Reactions
  12. Module 4: Drivers of Reactions
    4.1 Energy Changes in Chemical Reactions
  13. 4.2 Enthalpy and Hess's Law
  14. 4.3 Entropy and Gibbs Free Energy
  15. Module 5: Equilibrium and Acid Reactions
    5.1 Static and Dynamic Equilibrium
    5 Topics
  16. 5.2 Factors that Affect Equilibrium
    2 Topics
  17. 5.3 Calculating the Equilibrium Constant
    2 Topics
  18. 5.4 Solution Equilibria
  19. Module 6: Acid/Base Reactions
    6.1 Properties of Acids and Bases
    7 Topics
  20. 6.2 Using Brønsted–Lowry Theory
    2 Topics
  21. 6.3 Quantitative Analysis
    1 Topic
  22. Module 7: Organic Chemistry
    7.1 Nomenclature
    2 Topics
  23. 7.2 Hydrocarbons
    2 Topics
  24. 7.3 Products of Reactions Involving Hydrocarbons
  25. 7.4 Alcohols
    1 Topic
  26. 7.5 Reactions of Organic Acids and Bases
  27. 7.6 Polymers
    2 Topics
  28. Module 8: Applying Chemical Ideas
    8.1 Analysis of Inorganic Substances
    3 Topics
  29. 8.2 Analysis of Organic Substances
  30. 8.3 Chemical Synthesis and Design
  31. Working Scientifically
    Working Scientifically Overview
    1 Topic
Lesson 16, Topic 1
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Le Chatelier’s Principle

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Le Chatelier’s principle states:

When a change is made to an equilibrium system, the system moves to counteract the imposed change and restore the system to a new equilibrium.

Le Chatelier’s Principle

The new equilibrium position differs from the old equilibrium position. The position of an equilibrium refers to the comparative concentrations of reactants and products.

Homogenous Equilibrium = All products and reactants are of the same state, e.g. all gaseous or aqueous.

Heterogenous Equilibrium = Products and reactants are in multiple states


Take the following reaction as an example:

2\text{CrO}_4^{2-} (aq) + 2\text H^+ (aq) ⇌ \text {Cr}_2 \text 0_7^{2-} (aq) + \text H_2 \text O (l)

The chromate ion is yellow and the dichromate ion is orange.

Increasing the concentration of hydrogen ions is a change in the system. According to Le Chatelier’s principle, to counteract the change the hydrogen ion concentration is decreased by chromate ions reacting with some of the hydrogen ions to produce more dichromate ions. Thus the system becomes more orange. A new equilibrium position is established.

Gaseous Equilibria

In homogenous gaseous systems, changes in pressure (caused by changes in the volume of the vessel) will lead to changes in concentrations of reactants or products.

In general an increase in total pressure by volume reduction causes a gaseous equilibrium shift to the side of fewer gas molecules. If the number of particles is the same on each side of the equilibrium, then a change in total pressure has no effect.


Take the following reaction as an example:

\text N_2 \text O_2 (g) ⇌ 2\text {NO}_2 (g)

N₂O₂ is colourless and NO₂ is brown.

If the volume of the reaction vessel is increased, total pressure decreases (Boyle’s law). According to Le Chatelier’s principle, the system counteracts the change by shifting to the right to increase the number of particles. The system will become more brown.

Heterogenous Equilibria


A saturated sodium chloride solution is an example of a heterogenous equilibrium. Solid salt crystals are in equilibrium with ions in solution:

\text {NaCl} (s) ⇌ \text {Na}^+ (aq) + \text{Cl}^+ (aq)

If drops of concentrated hydrochloric acid are added to the saturated solution, the concentration of chloride ions increases. This disturbance shifts the equilibrium to the left to counteract the change. More NaCl crystals form.

If additional NaCl crystals are added to the saturated solution, no change is observed as the sodium and chloride ions are already at their maximum concentration in the water.

Adding water will dilute the concentration of the ions. The equilibrium shifts right as more salt crystals dissolve to restore the system to saturation.


The thermal decomposition of solid calcium carbonate in a closed system is an example of a heterogenous reaction because the carbon dioxide produced is gaseous:

\text {CaCO}_3 (s) ⇌ \text {CaO} (s) + \text {CO}_2 (g)

The equilibrium is endothermic in the forward direction. If the system is heated to a higher temperature, the equilibrium shifts to the right (according to Le Chatelier’s principle) to counteract the change and the partial pressure of carbon dioxide increases.

The addition of more solid calcium oxide or solid calcium carbonate causes no change to the equilibrium position.

The addition of more carbon dioxide gas raises the gas pressure and causes the equilibrium to shift to the left to counteract the change. The partial pressure of carbon dioxide will decrease until the new equilibrium is established.