Ch 6Equilibrium — Class 11 Chemistry, concept-first.

Objectives After studying this chapter, you should be able to: - identify the dynamic nature of equilibrium in physical processes (solid–liquid, liquid–vapour, solid–vapour, and dissolution of solids/…

Equilibrium isn’t something that happens only in a beaker of chemicals — it’s all around you. When you leave a glass of water open, the water level slowly drops; if you seal it, the level stays consta…

The simplest way to understand equilibrium in a chemical system is to look at a physical change: the coexistence of ice and water.

The equilibrium between a liquid and its vapour is a dynamic state that we encounter in everyday life — a covered water bottle, a sealed perfume vial, or steam above hot tea.

When a solid sublimes, it passes directly into the vapour phase without melting. If you place solid iodine in a closed vessel, you will soon see violet vapour filling the space above the solid.

Everyday experience tells us that only a fixed amount of a solid can dissolve in a given amount of liquid at a fixed temperature.

Before diving into the list of characteristics, it helps to understand what a physical process at equilibrium actually looks like. Think of a sealed bottle of water left on a table.

Chemical reactions, like physical processes, can reach a state of equilibrium. The key difference is that chemical equilibrium involves the transformation of substances into different chemical species…

When a reversible reaction reaches equilibrium, the mixture of reactants and products that remains is called an equilibrium mixture.

A chemical equilibrium is called homogeneous when all the reactants and products exist in the same physical phase.

So far, every equilibrium constant you have seen has been written using molar concentrations — the familiar square-bracket notation , with the constant called .

Equilibrium in a system that contains more than one phase is called heterogeneous equilibrium. Until now, we have mostly looked at reactions where everything is in the same phase — all gases, or all i…

Before we examine how equilibrium constants are used, we need to fix the key features that govern them.

The equilibrium constant (or ) is not just a number you calculate from concentrations at equilibrium — it tells you, at a glance, how far a reaction will go before it stops.

The equilibrium constant tells us where a reaction ends up — the ratio of products to reactants at equilibrium.

When you know the initial concentrations of reactants and products but have no direct measurement of what the system looks like at equilibrium, you need a systematic method to find the equilibrium con…

The equilibrium constant for a reaction is a number that tells you where equilibrium lies, but it is not a kinetic quantity — it does not depend on how fast the reaction gets there.

The central goal in chemical synthesis is to get the maximum possible conversion of reactants into products while using the least energy.

When a chemical system is at equilibrium, the forward and reverse reaction rates are equal, and the concentrations of reactants and products are constant.

When we talk about changing pressure in a gaseous reaction, we almost always mean changing the volume of the container.

When an inert gas — a gas that does not participate in the reaction — is added to a system at equilibrium while the volume is kept constant, the equilibrium position does not shift.

When you change the concentration, pressure, or volume of a system at equilibrium, the equilibrium constant itself does not change.

A catalyst is a substance that accelerates a chemical reaction without itself being consumed. In the context of equilibrium, its effect is often misunderstood.

You have already seen how changing the concentration of a reactant or product can shift the position of an equilibrium. One example you encountered involved ions:

Acids, bases and salts are everywhere around us. The hydrochloric acid in your stomach — about 1.2 to 1.5 litres secreted daily by the stomach lining — is essential for digestion.

The Arrhenius theory was the first successful attempt to define acids and bases in terms of their behaviour in water.

The older idea of acids and bases — that acids are substances that taste sour, turn blue litmus red, and liberate hydrogen gas with certain metals, while bases taste bitter, feel soapy, and turn red l…

In 1923, G.N. Lewis proposed a broader definition that shifted the focus from the proton to the electron pair.

The Arrhenius concept is most useful when we talk about acids and bases in water, because most chemical and biological ionizations happen in aqueous solution.

Water occupies a unique position in acid-base chemistry because it can act as both an acid and a base. This dual behaviour is called amphoteric nature.

The concentration of hydronium ions in a solution is often a very small number expressed in scientific notation. Working with such numbers directly can be cumbersome.

A weak acid, unlike a strong acid, does not completely dissociate into ions when dissolved in water. Instead, it establishes an equilibrium between the undissociated acid molecules and the ions produc…

A weak base, like a weak acid, does not fully dissociate in water. When a general weak base MOH is placed in water, an equilibrium is established between the unionized base and its ions:

The strength of an acid is measured by its acid dissociation constant , and the strength of a base by its base dissociation constant .

So far we have dealt with acids that donate a single proton () per molecule — monoprotic acids like or .

Having learnt to calculate the pH of acid solutions, a natural question follows: why does one acid donate its proton more readily than another? The extent of dissociation of an acid HA depends on two…

When a weak acid or weak base is placed in water, it establishes an equilibrium between the undissociated molecule and its ions.

When an acid and a base react in definite proportions, they form a salt. In water, these salts undergo ionization, releasing cations and anions.

Many fluids in the body — blood, urine, and others — have a very specific pH. A healthy person’s blood, for instance, stays close to pH 7.4.

The ability to prepare a buffer solution of a desired pH is a practical application of acid-base equilibrium.

The solubility of ionic solids in water spans an enormous range. Some salts, like calcium chloride, are so soluble they are hygroscopic — they absorb water vapour from the atmosphere.

When a sparingly soluble ionic solid like barium sulphate is placed in water, it does not dissolve completely.

When a sparingly soluble salt is in equilibrium with its saturated solution, the product of the concentrations of its ions (raised to appropriate powers) equals .

- Dynamic equilibrium: In a closed system, the forward and reverse reaction rates become equal — the system appears static but is microscopically active.

These are hands-on activities the NCERT textbook suggests so you can see equilibrium ideas from this unit at work outside the equations, using nothing more than pH paper, common kitchen/lab solutions,…

This chapter closes with the NCERT's own end-of-chapter exercise questions (Exercises 6.1–6.73), listed below with our own worked solutions.

Higher-order thinking problems from the NCERT Exemplar.