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 → Classification of elements and periodicity in properties → Periodic trends in properties ↓


Oxidation states


The concept of oxidation state is an essential part of understanding how elements react with each other. It is particularly important in the context of redox reactions, where there is a transfer of electrons. The oxidation state, also known as the oxidation number, helps us keep track of electrons during chemical reactions. This concept is important in explaining various chemical properties and predicting the outcomes of reactions.

What is the oxidation state?

The oxidation state is a number assigned to an element in a chemical compound indicating the number of electrons lost (or that can be considered lost) and gained (or that can be considered gained) by an atom of that element in the compound. The concept is used primarily to determine how electrons are transferred in chemical reactions, especially redox reactions. Generally, the oxidation state is positive for metals (indicating a loss of electrons) and negative for nonmetals (indicating a gain of electrons).

Assigning oxidation state

Determining the oxidation state involves some fundamental rules:

  • The oxidation state of an element in its pure form is always zero. For example, O_2 and N_2 both have oxidation state of 0.
  • For monatomic ions, the oxidation state is the same as the charge of the ion. For example, the oxidation state of Na+ is +1, and that of Cl- is -1.
  • In compounds, the oxidation state of hydrogen is usually +1, and the oxidation state of oxygen is usually -2. Exceptions exist, such as in H_2O_2 where the oxidation state of oxygen is -1.
  • In a neutral compound, the sum of the oxidation states must be zero. For example, in H2O, the sum of the oxidation states is (+1) * 2 + (-2) = 0.
  • For polyatomic ions, the sum of the oxidation states must equal the charge of the ion. For example, in SO42-, assuming that sulfur has an oxidation state of +6 and oxygen has an oxidation state of -2, the sum is 6 + 4*(-2) = -2.

Understanding through visual examples

Let us use some examples to better understand how oxidation states are determined.

Example 1: Water (H2O)

H - Oxidation state +1 2 x H = 2(+1) = +2 O - Oxidation state -2 Sum = +2 + (-2) = 0

Water is a neutral molecule, so the sum of the oxidation states is zero.

Example 2: Ammonium ion (NH4+)

N - Oxidation state -3 H - Oxidation state +1 4 x H = 4(+1) = +4 Total = -3 + 4 = +1

The total oxidation number is 1, which is equal to the charge on the ammonium ion.

Oxidation states of elements depending on location in the periodic table

Oxidation states are affected by the element's location in the periodic table, its group, and its period. Different groups of elements exhibit characteristic oxidation states:

  • Group 1 (Alkali Metals): The oxidation state of these elements is usually +1.
  • Group 2 (Alkaline earth metals): Here, the typical oxidation state is +2.
  • Group 17 (Halogens): These often display an oxidation state of -1. However, their oxidation state can become positive when combined with oxygen or other halogens.
  • Group 18 (Noble Gases): These elements have oxidation state of 0 because they are mostly inert and do not form compounds easily.

Oxidation states in a period

As you move from left to right across a period in the periodic table, oxidation states generally become more positive due to the increasing nuclear charge. Elements on the right side of a period gain electrons to achieve a full outer shell, which usually leads to negative oxidation states. Here is a simple representation:

Group 1 +1 Group 13 +3 Group 14 +4 Group 15 -3 +5 Group 16 -2 +6 Group 17 -1 +7

Oxidation states one group down

As you move down a group in the periodic table, the oxidation states of the elements generally remain the same. This is because each group has the same number of electrons in its outermost shell, which mainly determines their chemical reactivity and oxidation states.

Special cases

While general rules apply widely, some elements exhibit different oxidation states in different compounds due to special conditions or external factors. For example, transition metals often exhibit multiple oxidation states. Consider chromium and manganese, where their normal oxidation states can vary widely within a series of compounds:

  • Cr: usually displays +2, +3, and +6
  • Mn: can show +2, +4, +6, and even +7

The role of oxidation states in redox reactions

Redox reactions are the main processes of oxidation and reduction, which involve the transfer of electrons between chemical species. Understanding oxidation states is important for balancing redox reactions, because a change in oxidation state is always complemented by an equivalent change in the opposite direction. This means that the total increase in oxidation states must equal the total decrease.

Here's a simple example of a redox reaction:

Example: Oxidation of zinc with copper sulphate:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) Zn oxidation: 0 to +2 (loss of 2 electrons) Cu reduction: +2 to 0 (gain of 2 electrons)

Conclusion

The oxidation state is an important concept in chemistry that provides crucial insight into the behavior of elements during chemical reactions. The method of assigning oxidation states provides a structured way to understand electron transfer processes and makes balancing chemical equations more intuitive. By understanding oxidation states, chemists can predict reaction products, understand the electronic nature of bonds, and explore the rich variety of chemical reactions. With continued study, the application of oxidation states extends beyond basic chemistry, playing a fundamental role in advanced studies in inorganic, organic, and physical chemistry.


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Oxidation states

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