States of Matter & Chemical Equilibrium

7. States of Matter

Matter exists in various states, primarily solid, liquid, and gas, each characterized by distinct physical properties. Understanding these states is crucial for comprehending the behavior of substances.

7.1 Gaseous State

Gases are characterized by their ability to fill any container, low density, and high compressibility. Their behavior is largely explained by the Kinetic Theory of Gases.

7.1.1 Kinetic Theory of Gas and Its Postulates

The Kinetic Molecular Theory of Gases provides a microscopic model to explain the macroscopic properties of gases. Its key postulates are:

• Gases consist of a large number of identical particles (atoms or molecules) that are in constant, random motion.
• The volume occupied by the gas particles themselves is negligible compared to the total volume of the container.
• Intermolecular forces of attraction or repulsion between gas particles are negligible.
• Collisions between gas particles and with the walls of the container are perfectly elastic (no net loss of kinetic energy).
• The average kinetic energy of the gas particles is directly proportional to the absolute temperature (in Kelvin).

7.1.2 Gas Laws

The behavior of gases under varying conditions of pressure, volume, and temperature is described by several empirical laws.

7.1.2.1 Boyle’s Law and Charles' Law

• Boyle's Law: At constant temperature, the pressure of a fixed amount of gas is inversely proportional to its volume (P ∝ 1/V or PV = constant).
• Charles' Law: At constant pressure, the volume of a fixed amount of gas is directly proportional to its absolute temperature (V ∝ T or V/T = constant).

7.1.2.2 Avogadro's Law

At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of the gas (V ∝ n or V/n = constant). This implies that equal volumes of all gases, at the same temperature and pressure, contain the same number of molecules.

7.1.2.3 Combined Gas Equation

Combining Boyle's, Charles', and Gay-Lussac's laws, we get the combined gas law: (P1V1)/T1 = (P2V2)/T2. This equation is useful when none of the variables (pressure, volume, temperature) are constant.

7.1.2.4 Dalton's Law of Partial Pressure

For a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of individual gases (P_total = P1 + P2 + P3 + ...). The partial pressure of a gas is the pressure it would exert if it alone occupied the entire volume.

7.1.2.5 Graham's Law of Diffusion

The rate of diffusion or effusion of a gas is inversely proportional to the square root of its molar mass (Rate ∝ 1/√M). Lighter gases diffuse and effuse faster than heavier gases.

7.1.3 Ideal Gas and Ideal Gas Equation

An ideal gas is a hypothetical gas whose particles occupy negligible space and have no interactions. While no real gas is truly ideal, many gases behave ideally under conditions of high temperature and low pressure. The ideal gas equation combines all the gas laws into a single expression: PV = nRT.

7.1.4 Universal Gas Constant and Its Significance

The 'R' in the ideal gas equation (PV = nRT) is the universal gas constant. Its value depends on the units used for pressure, volume, and temperature. Its significance lies in linking the macroscopic properties of gases to the number of moles and temperature, serving as a fundamental constant in chemistry and physics.

7.1.5 Deviation of Real Gas from Ideality

Real gases deviate from ideal behavior at high pressures and low temperatures. This is because, under these conditions, the volume of gas molecules is no longer negligible, and intermolecular forces become significant. The Van der Waals equation is a modified ideal gas equation that accounts for these deviations.

7.2 Liquid State

Liquids have a definite volume but take the shape of their container. They exhibit stronger intermolecular forces than gases but weaker than solids, allowing molecules to move past each other.

7.2.1 Physical Properties of Liquids

7.2.1.1 Evaporation and Condensation

• Evaporation: The process where liquid molecules gain enough kinetic energy to escape from the liquid surface into the gas phase, occurring at any temperature.
• Condensation: The reverse process, where gas molecules lose energy and return to the liquid phase.

7.2.1.2 Vapour Pressure and Boiling Point

• Vapour Pressure: The pressure exerted by the vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system.
• Boiling Point: The temperature at which the vapor pressure of a liquid equals the external atmospheric pressure, allowing the liquid to rapidly turn into vapor throughout its bulk.

7.2.1.3 Surface Tension and Viscosity (Qualitative Idea Only)

• Surface Tension: The cohesive forces between liquid molecules cause the surface of a liquid to behave like a stretched elastic membrane, minimizing its surface area. This explains phenomena like water striders walking on water.
• Viscosity: A measure of a fluid's resistance to flow. It arises from the internal friction between layers of fluid molecules. High viscosity means the liquid flows slowly (e.g., honey), while low viscosity means it flows easily (e.g., water).

7.2.2 Liquid Crystals and Their Applications

Liquid crystals are states of matter that have properties between those of conventional liquids and solid crystals. They can flow like liquids but also exhibit long-range order in at least one dimension. They are widely used in Liquid Crystal Displays (LCDs) due to their ability to align in response to an electric field, modulating light transmission.

7.3 Solid State

Solids have a definite shape and volume, with particles tightly packed in fixed positions, exhibiting strong intermolecular forces and limited molecular motion (vibrations).

7.3.1 Types of Solids

Solids can be broadly classified based on the nature of their constituent particles and the forces holding them together:

• Ionic Solids: Held by electrostatic forces between ions (e.g., NaCl).
• Covalent Network Solids: Atoms held by covalent bonds in a continuous network (e.g., diamond, SiO2).
• Molecular Solids: Discrete molecules held by weak intermolecular forces (e.g., ice, sugar).
• Metallic Solids: Metal atoms held by metallic bonds in a 'sea of electrons' (e.g., copper, iron).

7.3.2 Amorphous and Crystalline Solids

• Crystalline Solids: Possess a highly ordered, repeating three-dimensional arrangement of particles (e.g., salt, quartz). They have sharp melting points.
• Amorphous Solids: Lack a long-range ordered structure; their particles are arranged randomly (e.g., glass, plastic). They soften gradually over a range of temperatures.

7.3.3 Efflorescent, Deliquescent and Hygroscopic Solids

• Efflorescent Solids: Crystalline solids that lose their water of crystallization when exposed to dry air (e.g., washing soda, Na2CO3·10H2O).
• Deliquescent Solids: Solids that absorb so much moisture from the air that they dissolve and form a solution (e.g., NaOH, CaCl2).
• Hygroscopic Solids: Solids that absorb moisture from the air without dissolving or changing their physical state (e.g., anhydrous CuSO4, silica gel).

7.3.4 Crystallization and Crystal Growth

Crystallization is the process by which solid crystals precipitate from a solution, melt, or more rarely deposit directly from a gas. It's a purification technique. Crystal growth refers to the subsequent increase in size of these formed crystals, influenced by factors like temperature, concentration, and presence of impurities.

7.3.5 Water of Crystallization

Water of crystallization refers to water molecules that are chemically bound within the crystal structure of a salt or other compound. These water molecules are part of the crystal lattice and give the crystal its characteristic shape and properties (e.g., CuSO4·5H2O, hydrated copper(II) sulfate).

7.3.6 Introduction to Unit Crystal Lattice and Unit Cell

• Crystal Lattice: A three-dimensional arrangement of points that represents the locations of atoms, ions, or molecules in a crystal. It is an imaginary framework.
• Unit Cell: The smallest repeating unit of a crystal lattice that, when repeated in three dimensions, generates the entire crystal. It defines the geometry and symmetry of the crystal.

8. Chemical Equilibrium

Chemical equilibrium is a state in reversible reactions where the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant.

8.1 Physical and Chemical Equilibrium

• Physical Equilibrium: Involves changes in physical states or properties, not chemical composition (e.g., liquid water ⇌ water vapor in a closed container, or ice ⇌ liquid water at 0°C).
• Chemical Equilibrium: Involves a balance between forward and reverse chemical reactions, where reactants are converted to products and vice versa (e.g., N2(g) + 3H2(g) ⇌ 2NH3(g)).

8.2 Dynamic Nature of Chemical Equilibrium

Equilibrium is not static; rather, it is dynamic. This means that at equilibrium, the forward and reverse reactions are still occurring, but at equal rates. So, there is a continuous interconversion of reactants and products, but with no net change in their concentrations.

8.3 Law of Mass Action

The Law of Mass Action states that the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants, each raised to a power equal to its stoichiometric coefficient in the balanced chemical equation.

8.4 Expression for Equilibrium Constant and Its Importance

For a general reversible reaction: aA + bB ⇌ cC + dD, the equilibrium constant (Kc for concentrations, Kp for partial pressures) is expressed as:

Kc = ([C]^c [D]^d) / ([A]^a [B]^b)

Importance: The value of the equilibrium constant indicates the extent to which a reaction proceeds towards products at equilibrium. A large K value means products are favored, while a small K value means reactants are favored.

8.5 Relationship between Kp and Kc

For reactions involving gases, the equilibrium constant can be expressed in terms of partial pressures (Kp) or molar concentrations (Kc). The relationship between them is given by:

Kp = Kc(RT)^Δn

Where R is the gas constant, T is the absolute temperature, and Δn is the change in the number of moles of gaseous products minus the number of moles of gaseous reactants (Δn = (c+d) - (a+b)).

8.6 Le Chatelier’s Principle

Le Chatelier's Principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. These conditions include changes in concentration, pressure (for gases), and temperature. While numerical problems are not required, understanding the qualitative effects is crucial:

• Effect of Concentration: Adding reactant shifts equilibrium to the right (produces more product); removing product shifts to the right.
• Effect of Pressure: Increasing pressure shifts equilibrium towards the side with fewer moles of gas; decreasing pressure shifts towards more moles of gas.
• Effect of Temperature: For an exothermic reaction, increasing temperature shifts equilibrium to the left (favors reactants); for an endothermic reaction, increasing temperature shifts to the right (favors products).
• Effect of Catalyst: A catalyst speeds up both forward and reverse reactions equally, so it helps reach equilibrium faster but does not change the position of equilibrium or the value of K.