Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → States of matter ↓


Ideal Gas Equation and Applications


The ideal gas equation is an essential concept in chemistry, providing a mathematical relationship that describes how an ideal gas behaves under various conditions. Let's begin our in-depth exploration of this concept and its various applications.

Understanding gases

To understand the ideal gas equation, it is important to understand the nature of gases. Gases are one of the elementary states of matter along with solids and liquids. They have unique characteristics, such as:

  • Gases assume the shape and volume of their container.
  • They are highly compressible compared to solids and liquids.
  • Gas molecules are in constant, random motion and expand rapidly to fill any available space.
  • They exert pressure equally in all directions as a result of molecular collisions with the walls of their container.

Ideal gas law

The ideal gas law is a fundamental equation that describes the behavior of ideal gases. It combines several simpler laws called Boyle's law, Charles' law, Avogadro's law, and Gay-Lussac's law. It is represented by the equation:

PV = nRT

Where:

  • P is the pressure of the gas
  • V is the volume of the gas
  • n is the number of moles of gas
  • R is the ideal gas constant
  • T is the temperature of the gas in Kelvin

Components of the ideal gas law

Pressure (P)

Pressure is the force exerted by gas molecules on the walls of their container. It can be measured in different units, such as atmospheres (atm), pascals (Pa), or millimeters of mercury (mmHg).

Volume (V)

Volume is the space the gas occupies. It is usually measured in liters (L) or cubic meters (m³).

Number of moles (n)

The number of moles represents the amount of gas molecules present. It is a measure of the amount of matter and is important in determining how many gas particles are in the system.

Temperature (T)

Temperature is a measure of the average kinetic energy of gas molecules and is usually measured in Kelvin for the purposes of the ideal gas law. To convert from Celsius to Kelvin, use the formula:

T(K) = T(°C) + 273.15

Ideal gas constant (R)

The ideal gas constant (R) is a proportionality constant. Its value depends on the units used for pressure, volume, and temperature. Common values include:

  • R = 8.314 text{ J/(mol·K)}
  • R = 0.0821 text{ L·atm/(mol·K)}

Applications of the ideal gas law

Calculating changes in volume and pressure of a gas

The ideal gas law allows us to calculate changes in pressure, volume, or temperature in a closed gas system. For example, if we know the initial conditions of a gas and how one of those variables changes, we can figure out how the other variables will be affected.

Example: Suppose you have 1 mol of gas at a pressure of 1 atm, a temperature of 273 K, and a volume of 22.4 liters. If the temperature is increased to 300 K, use the ideal gas law to find the new volume, assuming the pressure remains constant.

P1V1/T1 = P2V2/T2 (1 atm)(22.4 L)/(273 K) = (1 atm)(V2)/(300 K) V2 = (22.4 L)(300 K)/(273 K) ≈ 24.6 L

Determination of the molar mass of gases

The ideal gas law can also help determine the molar mass of an unknown gas, since it allows us to solve for n and use the total mass of a gas sample.

Example: A certain gas weighs 10 g, it fills a 5 liter container, exerts a pressure of 750 mmHg, and its temperature is 298 K. Calculate its molar mass.

PV = nRT n = PV/RT n = (750 mmHg)(5 L)/(62.36367 L·mmHg/(mol·K))(298 K) Molar Mass = mass/n Molar Mass = 10 g/n

Using conceptual diagrams

Here is a conceptual diagram showing the relationship between volume and temperature, often called Charles's law:

Temperature(K) Volume(L) v ∝ t

Real-life applications

Weather balloons

Weather balloons expand as they rise because the pressure decreases as you rise higher in the atmosphere. The ideal gas law helps predict how much the balloon will expand, which is important for making sure the balloon doesn't burst.

Respiratory function in medicine

The ideal gas law plays a vital role in understanding how gases used in medical applications, such as oxygen tanks, behave when subjected to varying pressures and temperatures. This understanding is essential to ensure patient safety and effective dose delivery.

Industrial applications

The production of ammonia in the Haber process depends on optimizing conditions of pressure, temperature, and volume, a process that can be modeled using concepts of the ideal gas law.

Deviations from ideal gas behaviour

While the ideal gas law is effective for many applications, it does not fully describe real gases because it assumes no interactions between gas molecules and an infinite container volume for the molecules. Gases at high pressure or low temperature may deviate from ideal behavior.

The Van der Waals equation is a more complicated equation that corrects for these interactions:

(P + n²a/V²)(V - nb) = nRT
  • a is responsible for the attractive forces between the molecules.
  • b corrects for the volume occupied by the gas molecules.

Understanding the ideal gas law and its applications increases our ability to analyze and predict the behavior of gases in various chemical contexts, thereby aiding scientific and industrial progress.


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