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 ↓


Balance


In grade 11 chemistry, one of the fundamental concepts you will encounter is the intricacies of balancing, especially when it comes to chemical equations. Balancing chemical equations is an important skill because it ensures the conservation of mass which is based on the law of conservation of matter. It states that during a simple chemical change, there is no detectable increase or decrease in the amount of matter.

To understand why balance is important, consider a chemical reaction equation such as the combustion of methane:

CH 4   O 2 → CO 2   H 2 O

To begin with, if you count the atoms of each element on both sides, they won't match. The equation should be balanced so that the amount of each atom is the same on both the reactant and product sides. This mirrors real-world chemical reactions, where matter is conserved.

Steps to balance a chemical equation

Balancing chemical equations may seem complicated at first glance, but it can be done systematically with the following steps:

  1. Write the unbalanced equation.
  2. Count the number of atoms of each element on both sides.
  3. Starting with the most complex molecule, add coefficients to balance the atoms.
  4. Keep repeating the counting and adjustments until all elements are balanced.
  5. Make sure that the coefficients are in the lowest possible ratio.

Example

Let's use a step-by-step approach to balance the equation for methane combustion.

CH 4   O 2 → CO 2   H 2 O
  1. Unbalanced Equations: 
    CH 4   O 2 → CO 2   H 2 O
  2. Count the atoms:

    Reactants: C=1, H=4, O=2

    Product: C=1, H=2, O=3

  3. Add up the coefficients:

    To balance the hydrogen, adjust the water:

    CH 4   O 2 → CO 2   2H 2 O
  4. Recalculation and balancing oxygen:

    Reactants: C=1, H=4, O=2

    Product: C=1, H=4, O=4

    CH 4   2O 2 → CO 2   2H 2 O
  5. Check the minimum whole number ratio:

    All coefficients are already the smallest whole numbers.

Visualizing equilibrium with chemical reactions

Let's illustrate the steps of equilibrium with a simple example using only hypothetical molecules. Visual aids can help reinforce the concept.

Consider a hypothetical reaction:

A 2   B <-> AB 2

Visual example - Balance

A2 B AB 2

Importance of balancing equations

So why is balance so important? Besides obeying the law of conservation of mass, there are some practical reasons for it:

  • Predictability: It allows chemists to predict the amount of products formed in a reaction.
  • Efficiency: Helps determine the ratio of reactants that will be maximally converted into products.
  • Safety: Knowing how much of each product is made can indicate potential hazards or reaction conditions.

Common challenges and tips

Balancing can be challenging at first, but with practice, it becomes natural. Here are common obstacles and tips for dealing with them:

  • Polyatomic ions: Consider them as complete units if they appear unchanged on both sides of the equation.
  • Fractional coefficients: If fractions come up, multiply the entire equation by the denominator to clear them.
  • Complex molecules: To simplify the process, start balancing the most complex molecules first.

Equilibrium in real-world reactions

Balancing is not just an academic exercise. Consider the industrial production of ammonia via the Haber process:

N 2   H 2 → NH 3

to balance:

  1. Write the unbalanced equation: 
    N 2   H 2 → NH 3
  2. Count the atoms:

    Reactants: N=2, H=2

    Product: N=1, H=3

  3. Add up the coefficients:

    Balance N, then H:

    N 2   3H 2 → 2NH 3
  4. Re-check and ensure the minimum whole number ratio.

Conclusion

Balancing chemical equations is an important skill in chemistry, enabling a deeper understanding of chemical reactions and the principles that govern them. With practice, the process becomes intuitive, paving the way for more advanced studies in chemistry and its applications in the real world.

Chemical equilibrium is a fundamental concept in chemistry that describes the state of a reversible chemical reaction where the rates of the forward and backward reactions are equal. In this situation, the concentrations of reactants and products remain constant over time. Chemical equilibrium is a dynamic process, meaning that reactions continue to occur, but there is no net change in the concentrations of reactants and products.


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Balance

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